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George Mason UniversityGeneral Chemistry 212

Chapter 14Main Group Element Patterns

AcknowledgementsCourse Text: Chemistry: the Molecular Nature of Matter and

Change, 7th edition, 2011, McGraw-HillMartin S. Silberberg & Patricia Amateis

The Chemistry 211/212 General Chemistry courses taught at George Mason are intended for those students enrolled in a science /engineering oriented curricula, with particular emphasis on chemistry, biochemistry, and biology The material on these slides is taken primarily from the course text but the instructor has modified, condensed, or otherwise reorganized selected material.Additional material from other sources may also be included. Interpretation of course material to clarify concepts and solutions to problems is the sole responsibility of this instructor.

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Main-Group Elements Chapter Overview

Application of bonding, structure, and reactivity to Main-Group Elements● Hydrogen● Period 2 Elements – Trends across Periodic Table● Group 1A – The Alkali Metals● Group 2A – The Alkaline Earth Elements● Group 3A – The Boron Family● Group 4A – The Carbon Family● Group 5A – The Nitrogen Family● Group 6A – The Oxygen Family● Group 7A – The Halogens● Group 8A – The Noble Gases

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Main-Group Elements In chemistry and atomic physics the periodic table divides the

elements into 4 groups Main Group Elements

● s-block: 1 (IA); 2 (IIA)● p-block: 13 (IIIA); 14 (IVA); 15 (VA)

16 (VIA); 17 (VIIA); 18 (VIIIA) Transition (d-block) Elements 3 4 5 6 7 8 9 10 11 12 (IIIB IVB VB VIB VIIB VIIIB VIIIB VIIIB IB IIB) Lanthanides (f-block Elements)

● Elements in Period 6, group 3 (IIIB) whosef-subshells are being filled

Actinides (f-block Elements)● Elements in Period 7, group 3 (IIIB) whose

f-subshellls are being filled)

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Main Group ElementsGroup→ 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18↓ Period IA IIA IIIB IVB VB VIB VIIB (VIIIB VIIIB VIIIB) IB 2B IIIA IVA VA VIA VIIA VIIIA

1

2

3

4

5

6

7

s block

p block

d-block(Transition Metals)

f-block - Lanthanoid (ide) series) f-block - Actinoid (ide) series)

Main Group Elements

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Main Group Elements Periodic Table Numbering System

Old Systems● Old IUPAC (Used in Europe)

Used Roman Numerals I, II, III, IV, V, VI, VII, VIII) and Letters (A &B) to indicate group (columns)

The numbers roughly indicated the highest oxidation state of the elements, thus similar chemical properties

The letters A and B were designated to the left (A) and right (B) part of the table

● CAS System (Used America) Similar to Old IUPAC except that the letter “A” referred

to the Main group Elements and the letter “B” referred to the Transition Elements

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Main Group Elemetnts Periodic Table Numbering System

Old systems confusing

● The use of the letters A & B in the old systems led to a lot of confusion

New IUPAC System (Universally used)

● Numbers the groups increasingly from 1 -18 left to right on the standard periodic table incorporating the 10 Transition Elements groups

● These group numbers correspond to the number of s, p, and d orbital electrons added since the last noble gas element (in column 18)

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Main Group Elements Main Group Elements

Elements that belong to the "s" and "p" blocks

Counting the columns (groups 1- 8) across the table(ignoring the transition elements) gives 8 element groups which match the filling of the eight spaces for electrons in the ns and np subshells, ns2np6

One good aspect about the 1 to 8 group numbering system is that the group number indicates the number of valence (outer) electrons for atoms in the main group elements

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Main Groupd Elemtns The lightest Main Group members are represented by

Helium, Lithium, Beryllium, Boron, Carbon, Nitrogen, Oxygen, And Fluorine

Main group elements (with some of the lighter transition metals) are the most abundant elements on the earth, in the solar system, and in the universeThey are sometimes called the representative elements

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Main-Group Elements Hydrogen (1s1)

90% of all atoms in Universe are Hydrogen atoms Single Electron; Small Size No perfectly suitable position in

the periodic table Depending on the property,

Hydrogen fits better in 1A, 4A, 7A

+1 oxidation State (grp 1A? Relatively High Ionization Potential (grp 7?) Forms diatomic molecule (H2 - grp 7?) Shares electrons (grp 4?) Half-filled valence shell; ionization energy; electron

affinity, electronegativity, and bond energies most similar to group 4

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Hydrogen Chemistry Hydrogen BondingDipole-Dipole force between Hydrogen (H) and small, highly

electronegative atoms with lone electron pair:Nitrogen (N); Oxygen (O); Fluorine (F)

Highly reactive, combining with nearly every element Ionic (salt like) hydrides

Group 1A & 2A metals2Li(s) + H2(g) 2LiH(s) Lithium HydrideCa(s) + H2(g) CaH2(s) Calcium Hydride

In H2O, H- is a strong base that pulls H+ from water Na+H-(s) + H2O Na+(aq) + OH-(aq) + H2(g)

Hydride ion is also a strong reducing agentTi4+Cl4(l) + 4LiH(s) = Tio(s) + 4LiCl(s) + 2H2(g)

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Covalent Hydrides Hydrogen reacts with nonmetals to form covalent hydrides

CH4 NH3 H2O HF

Conditions for forming Covalent Hydrides depend on the reactivity of the nonmetal – the more stable, the more temperature & pressure required for formation

Ex: Ammonia – 400oC & 250 atm

N2(g) + 3H2(g) 2NH3(g) Horxn = -91.8 kJ

At low temperatures (-196oC) Hydrogen combines readily with reactive Fluorine (F2)

F2(g) + H2(g) 2HF Horxn = -546 kJ

Catalyst

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Metallic (Interstitial) Hydrides Many Transition Elements form metallic (interstitial)

hydrides, where Hydrogen molecules (H2) and Hydrogen atoms (H) occupy the holes in the metal’s crystal structure.

These are not compounds, but rather gas-solid solutions

They lack a Stoichiometric formula because metal can incorporate a variable amount of hydrogen, depending upon temperature and pressure

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Trends Across Periodic Table

Electrons fill 1 ns – 3 np orbitals according to Pauli Exclusion Principle and Hund’s Rule

d orbitals in lower Periods can be used to accommodate additional oxidation states

Atomic size generally decreases

1st ionization potential increases

Electronegativity increases Metallic character

decreases with increasing nuclear charge

Reactivity highest at right & left sides, less in middle

Bonding - metallic covalent none (noble gas)

Continued on next Slide

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Trends Across Periodic Table

Bonding between each element and an active nonmetal changes from ionic to polar covalent

Bonding between each element and an active metal changes from metallic to polar covalent to ionic

Acid-Base behavior of common element oxide in water changes from basic to amphoteric (acts as acid or base (H2O) to acidic as bond between element and oxygen becomes more covalent

Reducing strength decreases through the metals

Oxidizing strength increases through the nonmetals

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Group 1A - Alkali Metals (ns1) Lithium (Li), Sodium (Na); Potassium (K); Rubidium (Rb);

Cesium (Cs); Francium (Fr) Single electron relatively far from nucleus weak metallic bonding - attraction between delocalized

electrons and metal-ion cores in crystalline structure Low melting points, soft consistency Reactive Metals Powerful reducing agents – lose 1 electron becoming 1+

cations, donating the electron to other elements ns1 configuration forms salts readily (+1 cations) Low Heat of Atomization (oHatom ) – Recall Lattice Energy)

Energy to convert solid into individual gaseous atomsoHatom (Li>Na>K>Rb>Cs)

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Group 1A - Alkali Metals (ns1) Low Ionization Energy (IE) – Each alkali element has

the largest size and the lowest IE in its Period Size of atom decreases considerably when valence

electron is lost Lattice Energy – The atomic radius increases as you

move down a group. Since the square of the distance is inversely proportional to the force of attraction, lattice energy decreases as the atomic radius increases

For a given anion, the Lattice Energy become smaller as the cation becomes larger

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Group 1A - Alkali Metals (ns1) Solubility – Despite strong ionic attractions, the Group

1A salts are water soluble – attraction between the ions and the polar Water molecule creates highly Exothermic Heat of Hydration (Hhydr)

Entropy – Entropy increases as ions disperse going into solution overcoming the high lattice energy

Magnitude of Hydration Energy decreases as ionic size increases

H = -Hhydr

(Li+ > Na+ > K+ > Rb+ > Cs+

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Group 1A - Alkali Metals (ns1) Anomalous Behavior of Lithium

Lithium ion (Li+) is small and highly positive

Dissociation of Lithium salts, such as LiF, Li2CO3, LiOH, and Li3PO4, in water is much more difficult than similar salts of sodium (Na) and Potassium (K)

Only member of Alkali group that forms simple Oxide and Nitride, Li2O & Li3N, on reaction with O2 & N2 in air

Only Lithium forms organo-metalic molecular compounds with hydrocarbon groups from organic Halides

2Li(s) + CH3CH2Cl(g) CH3CH2Li(s) + LiCl(s)

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Group 1A - Alkali Metals (ns1) Reactions & Compounds of Alkali Metals

Alkali metals reduce Hydrogen in Water to form Hydrogen gas

2E(s) + 2H2O 2E+ + 2OH-(aq) + H2(g)Where E = any alkali metal (Li, Na, K, Rb, Cs)

Reaction becomes more vigorous down group Alkali metals reduce oxygen, but product depends on

the metal4Li(s) + O2(g) 2Li2O(s) oxide K(s) + O2(g) KO2(s) superoxide Alkali metals reduce Hydrogen to form ionic hydrides

2E(s) + H2(g) 2EH(s)

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Group 1A - Alkali Metals (ns1) Reactions & Compounds of Alkali Metals

Alkali metals (E) reduce Halogens (X) to form Halides2E(s) + X2 2EX(s) X = F, Cl, Br, I)

Sodium Metal (Na) can be produced from Molten NaCL and electricity

2NaCl(l) 2Na(l) + Cl2(g) Sodium Hydroxide (Lye) can be produced from Salt (NaCl),

water (H2O) and electrolysis2NaCl(s) + H2O(l) 2NaOH(aq) + H2(g) + Cl2(g)

In an ion-exchange process, water can be “softened” by removal of dissolved hard-water cations to displace Na+ ions from a “resin”

M2+(aq) + Na2Z(s) MZ(s) + 2Na+(aq)

(M = Mg, Ca: Z = resin)

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Group 1A - Alkali Metals (ns1)

atomic properties

physical properties

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Group 2A - Alkaline Earth Metals (ns2) Be, Mg, Ca, Sr, Ba, Ra (E2+ ions) Oxides (except Be) give basic (alkaline) solutions:

Ca(OH)2, Mg(OH)2

High melting points (higher lattice energy than 1A) Atomic & Ionic sizes

Smaller radii and higher ionization energy Increase in size down the group Combination of size, extra electron, and metallic

bonding result in stronger attractions between delocalized electrons and the atom cores

Thus, Melting Points and Boiling Points are much higher than 1A alkali metals

Harder & more dense than Alkali metals, but soft and lightweight compared to transition metals (Fe, Cr, etc)

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Group 2A - Alkaline Earth Metals (ns2) Even though the Alkaline Earth metals have higher ionization

potential, they still form ionic compounds (E2+), but Beryllium (Be) is an exception forming covalent bonds

Like Alkali metals, Alkaline Earth metals are strong reducing agents

Group 2A (Alkaline Earth) elements are reactive because the higher lattice energy of their compounds more than compensates for the large total Ionization Energy (IE) needed to form the 2+ cations

The higher Lattice Energy (from the smaller cation size) and higher Charge Density results in lower solubility

Ion-Dipole attraction is so strong that many slightly soluble 2A salts crystallize as “Hydrates”

Epsom salt – MgSO47H2O Gypsum – CaSO42H2O

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Group 2A - Alkaline Earth Metals (ns2) The anomalous behavior of Beryllium

Beryllium has smallest size; highest Ionization energy, and highest Electronegativity of the Alkaline Earth elements

Combined with the high charge density of the ion (Be2+) it polarizes the nearby electron clouds very strongly and causes extensive orbital overlap; this results in covalent bonding

BeF2 is the most ionic of the Beryllium compounds, but its melting point and electrical conductivity are relatively low compared to the other alkaline earth Fluorides

Unlike the other Alkaline Earth Metals, whose oxides are basic, BeO is amphoteric and does not react with water to form OH- ions

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Group 2A - Alkaline Earth Metals (ns2) Diagonal relationships: Lithium and Magnesium Certain Period 2 elements exhibit behaviors that are very

similar to those of the Period 3 elements immediately below and to the right

Lithium and Magnesium reflect similar atomic and ionic size Both elements form:

Nitrides, Hydroxides and Carbonates (CO3) that decompose with heat, Organic compounds with polar covalent metal-carbon bonds Salts with similar solubilities

3 relationships1. Li, Mg2. Be, Al3. B, Si

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Group 2A - Alkaline Earth Metals (ns2) Reactions & Compounds (E = Mg, Ca, Sr, Ba)

Metals reduce Oxygen (O2) to form Oxides

2E(s) + O2(g) 2EO(s)

Ba + O2 BaO2 (Barium Peroxide)

Larger metals reduce water to form hydrogen gas

E(s) + 2H2O(l) E2+aq) + 2OH- (aq) + H2(g)

Metals reduce Halogens to form ionic halides

E(s) + X2 EX2(s) (X = F, Cl, Br, I)

Most metals (Be exception) reduce Hydrogen to form ionic hydrides

E(s) + H2(g) EH2 (s) (except Be)

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Group 2A - Alkaline Earth Metals (ns2) Reactions & Compounds (E = Ca, Mg, Sr, Ba)

Most elements reduce Nitrogen to form ionic Nitrides3E(s) + N2(g) E3N2(s) (except Be)

Element Oxides are Basic (except for amphoteric BeO)EO(s) + H2O(l) E2+(aq) + 2OH-(aq)

All Carbonates undergo thermal decomposition to the oxide

ECO3(s) EO(s) + CO2(g) (CaO – Lime) Beryl (Be3Al2Si6O18) - Gemstone, source of Be Magnesium oxide (MgO) – Refractory material for

furnace bricks Alkyl Magnesium Halides – RMgX (R=Hydrocarbon)

Grignard Reagents – organic compound synthesis

heat

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Group 2A - Alkaline Earth Metals (ns2)

atomic properties

physical properties

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Group 3A – Boron Family (ns2np1)B Al Ga In Tl

Boron heads family, but other elements in group 3A exhibit diverse properties

Boron & Aluminum, especially Aluminum, are much more abundant than the others, but still quite rare

Group 3A elements include “p” orbitals for first time In Period 4 (transition elements) the “d” orbitals are

present Physical Properties are influenced by type of bonding

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Group 3A – Boron Family (ns2np1) Boron is a network covalent metalloid - Black, hard,

very high melting point A network solid or covalent network solid is a

chemical compound in which the atoms are bonded by covalent bonds in a continuous network

In a network solid there are no individual molecules and the entire crystal may be considered a macromolecule

Boron (metalloid) is much less reactive than the others members of the 3A group because it forms covalent bonds

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Group 3A – Boron Family (ns2np1) Other group members are metals – shiny, relatively soft

with low melting points Aluminum is more ionic; its low density and 3 valence

electrons make it a good electrical conductor Although Aluminum is a metal, its halides exist in the

gaseous state as covalent dimers - AL2Cl6 (contrast salts of group 1 & 2 metals)

Aluminum Oxide, Al2O3, is amphoteric (can act as an acid or base) rather than basic like the Group 1A & 2A metals

Although the other Group 3A elements are basically ionic they exhibit more Covalent character than similar 2A compounds.

3A cations are smaller with more charge density than 2A cations and they polarize an anion’s electron cloud more effectively

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Group 3A – Boron Family (ns2np1) Oxidation-Reduction (REDOX) behavior in Group 3A

Presence of Multiple Oxidation States

● In Groups 3A – 6A many of the larger elements (down the group) exhibit an oxidation state “two lower” than the A-Group number

● This lower state occurs when the atoms lose their np electrons, not the ns electrons.

● The lower oxidation state is the result of lower bond energies

● Bond energies decrease as the size of the atom and the bond length increase for elements lower in the group

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Group 3A – Boron Family (ns2np1) Increasing prominence of the low oxidation state

● When a group exhibits more than one oxidation state, the lower state becomes more prominent going down the Group

● All members of the 3A group exhibit the +3 state, but the +1 state appears first with some compounds of Gallium (Period 4)

● The +1 state becomes the most important state of Thallium (Period 6)

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Group 3A – Boron Family (ns2np1) Relative Basicity of Group 3 oxides

● Recall: A1 oxides (ionic charge +1 and more metallic) are more basic than A2 oxides (ionic charge +2 and less metallic)

● In general, oxides with the element in a lower oxidation state (less positive) are more basic than oxides with the element in a higher oxidation state

● For Indium oxides in Group A3, In+12O acts more like

a metal and is more basic than In+32O3

● The lower charge density of In+1 does not polarize the O-2 ion as much as the In+3 ion

● Thus, in E2O compounds, the E-O bonding is more ionic than in E2O3 compounds, thus; the O-2 ion is more available to act as a base – donate electron pair or accept a proton

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Group 3A – Boron Family (ns2np1) Boron Chemistry

Boron compounds are covalent (unique within group) Forms network covalent compounds or large molecules

with metals, H, O, N Electron deficient; uses two approaches to complete

octet ● Accepting a Bonding Pair from Electron-Rich atom

BF3(g) + NH3(g) F3B-NH3(g) (BF3 acts as acts as Lewis acid in accepting the electron

pair from the Nitrogen in NH3)

B(OH)3 + H2O(l) B(OH)4-(aq) + H+(aq)

(Acts as acid by accepting electron pair from H2O)Note: Water is acting as the base

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Group 3A – Boron Family (ns2np1) Boron Chemistry

Two approaches to filling octet (con’t)● Accepting electron pair from Electron-Rich atom

(con’t)Boron-Nitrogen compounds are similar in

structure to elemental Carbon and some of its organic compounds

Size, Ionization Energy, Electronegativity of Carbon is between Boron & Nitrogen

Ethane & Amine – Borane have the same number & electron configuration

C – C

B – N

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Group 3A – Boron Family (ns2np1) Boron Chemistry

Two approaches to filling octet (con’t)● Forming Bridge Bonds with Electron-Poor Atoms

Boron Hydrides - Boranes 2 types of B – H bonds

Normal electron-pair bondo sp3 orbital of B overlaps 1s orbital of H in each

of the four terminal B-H bonds

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Group 3A – Boron Family (ns2np1) Hydride Bridge Bond (3-center, 2 electron

bond)o Each B – H – B grouping is held together by only

two electronso Two sp3 orbitals, one from each B, overlap an H

1s orbital between themo Two electrons move through this extended

bonding orbital – one from one of the B atoms and the other form the H atom – and join the 2 B atoms via the H atom bridge

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Group 3A – Boron Family (ns2np1)

atomic properties

physical properties

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Group 3A – Boron Family (ns2np1) Reactions & Compounds

Elements react sluggishly, if at all, with water (H2O)2 Ga(s) + 6H2O(hot) 2Ga3+(aq) 6OH-(aq) + 3H2(g)

2Tl(s) + 2H2O(steam) 2Tl+(aq) +2OH-(ag) + H2(g)Note different oxidation numbers for Ga3+ & Tl+

All members form oxides when heated in pure O2

4E(s) + 3O2(g) 2E2O3(s) (E = B, Al, Ga, In)4Tl(s) + O2 2Tl2O3(s)

Oxide acidity decreases down the group:B2O3 > Al2O3 > Ga2O3 > In2O3 > TlO2

(weakly acidic) (strongly basic)

The +1 oxide (TlO2) is more basic than the +3 oxide

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Group 3A – Boron Family (ns2np1) Reactions & Compounds

All members reduce Halogens2E(s) + 3X2 2EX3 (E = B, Al, Ga, In)

2Tl(s) + X2 2TlX(s) Trihalides of AL, Ga, In are mostly ionic but

exist as dimers in the gas phase

Acid (H2SO4) treatment of Al2O3 produces Al2SO4, a colloid (coagulant) used in water purification

Al2O3 + 3H2SO4 Al2SO4(s) + 3H2O(l)

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Group 4A – Carbon Family (ns2np2) The whole range of elemental behavior occurs within the

4A group Non metalic Carbon (C) Metalloids (Silicon (Si) & Germanium (Ge) Metallic (Tin (Sn) & Lead (Pb) Newly synthesized element at bottom of group

Carbon forms the basis of “Organic Chemistry”

20,000,000 compounds Polymer Chemistry Biochemistry based on Carbon Geochemistry Electronic technologies bases on Si

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Group 4A – Carbon Family (ns2np2) Bonding effects on Physical Properties

Silicon has a much lower melting point than Carbon because of the longer, weaker bonds.

The melting point difference between Germanium (Ge) and Tin (sn) is due to the change from network covalent to metallic

Going from Group 3 to group 4 there are large increases in melting point and the Hfus because of the change from metallic to network covalent bonding

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Group 4A – Carbon Family (ns2np2) Allotropism: Different Forms of an Element

Elemental Carbon – Graphite & Diamond Different crystalline & molecular forms with different

physical properties● Carbon Allotropes

Graphite – Black, “greasy”, soft, more stable than diamond

Diamond – Colorless, electrical insulator, extremely hard

Bucky-Balls (Buckminsterfullerene) – soccer ball-shaped with the formula C60

● Tin Allotropes-tin – stable at room temperature & above-tin – stable below 13oC

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Group 4A – Carbon Family (ns2np2) Bonding Changes in Group 4A

Carbon – Covalent (intermediate EN)

Si & Ge – strong polar bonds (silicate minerals)

Tin (Sb) & Lead(Pb) – Metallic (Ionic)

Multiple Oxidation States

Carbon (+4)

Silicon (+4 more stable than +2)

Lead (+2 more stable than +4)

Elements with lower oxidation states act more like metals (more basic)

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Group 4A – Carbon Family (ns2np2) Highlights of Carbon Chemistry

Carbon, like other elements in “Period 2, is the anomalous element in the group

Carbon forms bonds with:

● Smaller Group 1A & 2A metals

● Many transition metals

● Halogens

● Neighbors B, Si, N, O, P, S

● Exhibits all possible oxidation states from +4 in CO2, and Halides to -4 in CH4

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Group 4A – Carbon Family (ns2np2) Highlights of Carbon Chemistry

Two main features of Carbon Chemistry● Catenation and the ability of Carbon to form

multiple bondsCarbon can form chains, branches, and rings

(aromatic & aliphatic)Multiple bonds – sigma (), Pi (), Triple ()The C-C bond is short enough for side-to-side

overlap of two half-filled 2p orbitals to form bonds that give rise to many diverse structures and reactivities of organic compounds

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Group 4A – Carbon Family (ns2np2) The Other 4A elements

E-E bonds become longer going down the group, with decreasing bond strength

C – C > Si – Si > Ge – Ge

The empty d shell orbitals make these chains susceptible to chemical attack – they are reactive

The long bonds are not suitable for overlap of p orbitals; thus, no bonds

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Group 4A – Carbon Family (ns2np2) Carbonates

Metal Carbonates are the main mineral

Marble, Limestone, Chalk, Coral, others Antacids – Calcium Carbonate & Stomach Acid

CaCO3(s) + 2HCl(aq) CaCl2(aq) + CO2(g) + H2O(l)

Limestone (CaCO3) deposits help moderate the effects of acid rain (H2SO4 & HNO3)

Carbon Dioxide (CO2)

Essential to all life as primary source of carbon in plants & animals through photosynthsis

Atmospheric buildup from motor vehicles and fossil fuel powerplant severely affect global climate

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Group 4A – Carbon Family (ns2np2) Silicon Chemistry

Silicon Halides are more reactive than Carbon Halides because Si (3s, 3p, 3d orbitals) has empty 3d orbitals available for bond formation

The Si – X bond is long but stronger than corresponding C – X bond

Si – X bond has some double bond character because of the presence of a bond and a different type of bond called a p,d- (side-to-side overlap of the Si d orbital and a Halogen p orbital

The impact of p,d- bonding on the structure of trisilylamine

Trisilylamine(SiH3)3N

trigonal planar

Trimethylamine

(CH3)3N

trigonal pyramidal

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Group 4A – Carbon Family (ns2np2) Silicon chemistry is dominated by the Silicon-Oxygen –

(Si–O) bond C – C bonds can repeat endlessly; similarly, the Si–O

bonds can also repeat forming large chains in Silicate minerals in the earths crust and Silicones, which are synthetic polymers with a large number of industrial applications

Silicate Minerals From common sand (SiO2) and clay to semiprecious

amethyst, Silicate minerals are the dominant form of matter on the earth

Oxygen is the most common element on earth and Silicon is the next most abundant

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Group 4A – Carbon Family (ns2np2) The Orthosilicate (–SiO4 –) grouping is the building unit for

Silicate minerals Zircon ZrSiO4 1 Unit Hemimorphite 2 Units [Zn4(OH2Si2O7H2O] Beryl 6 Units [Be3Al2Si6O18]

Silicon Polymers Manufactured Substances Alternating Si & O atoms with two Organic groups

bonded to each Silicon atom

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Group 4A – Carbon Family (ns2np2)

atomic properties

physical properties

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Group 4A – Carbon Family (ns2np2) Important Reactions

Group 4A elements are Oxidized by Halogens

E(s) + 2X2 EX4 (E = C, Si, Ge)The +2 Halides are more stable for Tin (Sb) & Lead (Pb)

SnX2 PbX2

The Elements are oxidized by Oxygen (O2)

E(s) + O2(g) EO2 (E = C, Si, Ge, Sn)The oxides are more basic (metallic) going down the group

Lead (Pb) forms the +2 oxide (PbO) - basic

In natural streams, Carbon Dioxide (CO2) forms a weakly “acidic” solution

CO2 + H2O H⇄ 2CO3(aq) H⇄ +(aq) + HCO3-(aq)

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Group 4A – Carbon Family (ns2np2) Air & Steam passed over hot coke (carbon) produce

gaseous fuel mixtures – producer gas & water gasC(s) + H2O(g) + air(g) CO(g) + CO2(g) + N2(g) + H2(g)

Note: This industrial reaction cannot be balanced Hydrocarbons (C & H only) react with Oxygen (O2) to form

CO2 & Water (H2O), a source of heat to yield steam (H2O) for electrical generation

CH4(g) + 2O2(g) CO2(g) + 2H2O(g) + Heat Certain metal Carbides react with water to produce

Acetylene (H-CC-H), used in oxyacetylene torchesCaC2(s) + 2H2O(g) Ca(OH)2(aq) + C2H2(g)

Acetylene is source material for organic compound synthesis and a fuel for Welding

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Group 4A – Carbon Family (ns2np2) Freon (chlorofluorocarbon) is formed from fluorinating

Carbon Tetrachloride

CCl4(l) + HF(g) CFCl3(g) + HCL(g)

Production of Trichlorofluoromethane (Freon-11) is being discontinued because it is an atmospheric pollutant

Silica (SiO2)is reduced to form elemental Silicon used in the manufacture of computer chips

SiO2(s) + 2C(s) Si(s) + CO2(g)

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Group 5A – Nitrogen Family (ns2np3) Widest range of physical behavior in 1st 5 groups Compounds of Nitrogen (gaseous nonmetal) and

Phosphorus (solid nonmetal) are important in industrial and environmental processes

Arsenic (As) and Antimony (Sb) are network covalent metalloids with highest melting points in group

Bismuth (Bi) exhibits metallic bonding Nitrogen (N) exists as diatomic molecules, which interact

through very weak dispersion force producing a boiling point 200 oC below room temperature

Phosphorus (P) is heavier and more polarizable than Nitrogen with stronger dispersion forces – higher melting point 44oC

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Group 5A – Nitrogen Family (ns2np3)

atomic properties

physical properties

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Group 5A – Nitrogen Family (ns2np3) Chemical Behavior in Group 5A Patterns

Nitrogen forms a maximum of 4 covalent bonds The other elements in the group can expand the

valence shell by using empty ‘d’ orbitals The noble gas configuration is attained by group 5

elements gaining 3 electrons – the first is Exothermic and the last two are Endothermic (requiring input of energy from surroundings

As in groups 3A & 4A, fewer oxidation states occur moving down the group with the lower oxidation state becoming prominent

Oxidation states● Nitrogen – from +5 to -3● P, As, Sb – +5 & +3● Bi – +3

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Group 5A – Nitrogen Family (ns2np3) Nitrogen Chemistry

Nitrogen Oxides – 6 stable forms

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Group 5A – Nitrogen Family (ns2np3) Nitrogen Oxoacids & Oxoanions

Oxoacid

Oxoanion

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Group 5A – Nitrogen Family (ns2np3) Nitric Acid (oxidizing agent) Reactions

Active Metal (Al in dilute HNO3 solution)8Al(s) + 30 HNO3(aq) 8Al(NO3)3 + 3NH4NO3 + 9H2O(l)

Al(s) + 30H+(aq) + 3(N+5O3)-(aq) 8Al+3(aq) + 3(N+3H4)+(aq) + 9H2O(l)

(net ionic equation) Less reactive metal (Cu), more conc HNO3 (N2+ forms)

3Cu(s) + 8HNO3(aq) 3Cu(NO3)2 + 4H2O(l) + 2N2+O(g)

3 Cu(s) + 8H+(aq) + 2(N+5O3)- 3Cu2+(aq) + 4H2O(l) + 2N+2O(g)

(net ionic equation) Copper with still more concentrated HNO3 (N4+ forms)Cu(s) + 4HN+5O3(aq) Cu+2(NO3)2(aq) + 2H2O(l) + 2N4+O2(g)Cu(s) + 4H+(aq) + 2NO3

-(aq) Cu+2(aq) + 2H2O(l) + 2NO2(g)

(net ionic equation)

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Group 5A – Nitrogen Family (ns2np3) Nitrates form when Nitric Acid (HNO3) reacts with the

hydroxides, oxides, or carbonates of metalsHNO3 + NaOH NaNO3 + H2O

Nitrous Acid (HNO2), a much weaker acid, is formed from the reaction of a strong acid (HCl) and metal Nitrites

NaNO2(aq) + HCl(aq) HNO2(aq) + NaCl(aq) Strong acid vs Weak acid – The more oxygen atoms

bonded to the central nonmetal, the stronger the acid Oxygen atom pulls electron density from the Nitrogen

atom, which in turn pulls electron density from the Oxygen of the O-H bond, facilitating the release of the H+ ion – The more Protons (H+) in solution, the stronger the acid.

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Group 5A – Nitrogen Family (ns2np3) Phosphorus Chemistry

Phosphorus forms two important Oxides● Tetraphosphorus Hexaoxide, P4O6

P (+3) has tetrahedral orientation with an Oxygen between pair of P atoms

Reacts with water to form Phosphorus acid, H3PO3

P4(s) + 3O2(g) P4O6(s) + 6H2O(l) 4H3PO3(l)Only two of the H atoms are acidic, the third is

bonded to the central P and does not dissociateDissociation is complete in strong base solution

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Group 5A – Nitrogen Family (ns2np3)● Tetraphosphorus Decaoxide (con’t)

Phosphorus (+5) oxidation stateP4(s) + excess 5O2 P4O10(s)P4O10(s) + 6H2O 4H3PO4(l)

● Phosphoric Acid (H3PO4) is a weak triprotic acidIn water it loses one proton to form H2PO4

-

In excess strong base all three protons dissociate to form the Phosphate ion, PO4

3- + 3H+

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Group 5A – Nitrogen Family (ns2np3) Diphosphate & Polyphosphates

Polyphosphates are formed by heating Hydrogen Phosphates (ex. Na2HPO4)

2Na2HPO4(s) Na4P2O7(s) + H2O(g) The Diphosphate ion, P2O7

4-, is the smallest of the polyphosphates consisting of tetrahedral PO4 units linked through a common Oxygen

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Group 5A – Nitrogen Family (ns2np3) Reactions & Compounds

Converting Nitrogen to other forms (fixing) is quite difficult because of the strength of the triple bond (NN) between the Nitrogen atoms. It can be fixed industrially by the Haber process

N2(g) + 3H2(g) 2NH⇄ 3(g) Non Nitrogen Hydrides from metal Phosphides

Ca3P2(s) + 6H2O(l) 2PH3(g) + 3Ca(OH)2

Halides formed by direct combination of elements

2E(s) + 3X2 2EX3 (E= P, As, Sb not N)

EX3 + X2 EX5 (all except N & Bi with X = F & Cl) P4 in basic solution increases & decreases oxidation

number

P4(s) + 3OH-(aq) + 3H2O P3+H3(g) + 3H2P1+O2-

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Group 6 – Oxygen Family (ns2np4) The Oxygen Family

First two members of group – gaseous nonmetallic oxygen (O) & solid nonmetallic sulfur (S) are among most important elements in industry, the environment and living organisms

Selenium (Se) & Tellurium (Te) are metalloids

Polonium (Po) is radioactive and only metal in the group

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Group 6 – Oxygen Family (ns2np4)

atomic properties

physical properties

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Group 6 – Oxygen Family (ns2np4) Oxygen Family vs Nitrogen Family

Groups 5A & 6A have similar Physical & Chemical Properties

Oxygen & Nitrogen are low-boiling Diatomic gases Phosphorus (P) & Sulfur (S) occur as polyatomic

molecules – P4 & S8

Arsenic (Ar) & Selenium (Se) occur as gray metalloids Antimony (Sb) & Tellurium more more metallic than

preceding group members, but display network convalent bonding

Bismuth & Polonium are metallic crystals Electrical conductivity increases down group as

bonding changes from individual molecules (insulators) to metalloid networks to metallic solids (conductors)

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Group 6 – Oxygen Family (ns2np4) Oxygen Family vs Nitrogen Family

Allotropism (two or more crystalline or molecular forms of an element) is more common in Group 6A than in Group 5A● Oxygen has 2 allotropes

Dioxygen O2 & Ozone O3

● Oxygen (O2) gas is colorless, odorless, paramagnetic (unpaired electrons attracted by outside magnetic field), and thermally stable

● Ozone (O3) gas is bluish, has pungent odor, is diamagnetic (paired electrons not affected by external magnetic field), decomposes in heat and Ultraviolet light.

● Ozone in upper atmosphere protects living organisms from Ultraviolet radiation

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Group 6 – Oxygen Family (ns2np4) Oxygen Family vs Nitrogen Family

Sulfur Allotropes● 10 forms ● Sulfur bonds to other Sulfur atoms

creating rings and chains● The most stable form is orthorhombic,

S8, a crown-shaped ring of 8 atoms

top view

side view

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Group 6 – Oxygen Family (ns2np4) Oxygen Chemistry vs Nitrogen Chemistry

Oxygen & Sulfur occur as anions more often than Nitrogen & Phosphorus

Oxygen & Sulfur bond covalently with almost all nonmetals

Selenium & Tellurium do some covalent bonding, whereas Polonium behaves like a metal

Oxygen has few oxidation states (O2- most common) The other elements in the family exhibit +6. +4,

-2 oxidation states, with the +4 state most common in Tellurium and Polonium

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Group 6 – Oxygen Family (ns2np4) Oxygen Chemistry vs Nitrogen Chemistry

Oxidizing strength of Oxygen is 2nd only to Fluorine

The other members of the group behave very little like oxygen being less electronegative and forming anions less often

Except for Oxygen, all elements of group 6A form foul smelling, poisonous, gaseous hydrides (H2E) upon treatment of the metal Sulfide, Selenide, etc., with an acid

FeSe(s) + HCl (aq) H2Se(g) + FeCl2(aq)

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Group 6 – Oxygen Family (ns2np4) Oxygen Chemistry vs Nitrogen Chemistry

Bonding & Thermal stability of Group 6A elements have several features in common● Only Water (H2O) forms Hydrogen bonds, so it melts

& boils at higher temperatures than the group 6A H2E compounds (E = S, Se, Te, Po)

● Bond angles drop from the nearly tetrahedral for H2O (104.5o) to around 90o for the Group 6A element hydrides (unhybridized p orbitals)

● E-H bond length increases (bond energy decreases) down group

● Thus, H2Te is stable above 0oC, but H2Po is only stable at extremely cold temperatures; it even decomposes from the heat generated by the radioactivity of the Polonium.

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Group 6 – Oxygen Family (ns2np4) Bonding & Thermal Stability (con’t)

● Except for Oxygen, Group 6A elements form a wide range of Halides

● The Halide structure and reactivity patterns depend on the sizes of the central atom and the surrounding Halogens

Radius of S < Se < Te < Po Sulfur – Many Fluorides, Few Chlorides, one

Bromide As the central atom becomes larger, the Halides

become more stable Tetrachlorides & Tetrabromides of Se, Te, Po

are known Tetraiodides of Te & Po are known Hexafluorides of S, Se, Te are known

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Group 6 – Oxygen Family (ns2np4)● Halide Structure (con’t)

Inverse relationship between bond length & bond strength does not explain pattern of Group 6 Halide formation

Crowding of lone electron pairs and Halogen (X) atoms around the central atom

With S (small central atom) the larger Halides further down group 7 would be too crowded, which explains why Sulfur Iodides don’t occur

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Group 6 – Oxygen Family (ns2np4)● Halide activity (con’t)

Sulfur Tetrafluoride vs Sulfur HexafluorideSF4 has a lone pair of unshared electrons and

empty d orbitals which can be involved in bonding – Highly reactive

SF6 uses all of the bonds allowable for S and is tightly packed – chemically inert

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Group 6 – Oxygen Family (ns2np4) Highlights of Oxygen Chemistry

Most abundant element of the Earth’s surface Many oxides – water, silicates, carbonates, phosphates Most free Oxygen (O2) has biological origin from

photosynthesis in algae and multicellular plants Although much more complicated, the basic reaction

between oxygen, CO2 and light to form carbohydrates can be represented as:

nH2O(l) + nCO2(g) nO2(g) + (CH2O)n

The reverse process of combustion and respiration produce CO2

Every element, except He, Ne, Ar (noble gases) form at least one oxide

Some oxides have Endothermic heats of reaction, while others have Exothermic ones

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Group 6 – Oxygen Family (ns2np4) Highlights of Sulfur Chemistry

Two important oxides – SO2 & SO3

● Sulfur Dioxide (+4 oxidation state) is a colorless choking gas that forms when S, H2S, or a metal sulfide burns in air (oxygen)

2H2S (g) + 3O2(g) 2H2O(g) + 2SO2(g)FeS2(s) + 11O2(g) 2Fe2O3(s) + 8SO2(g)

● Sulfur Dioxide dissolves in water to form Sulfurous acid (H2SO3) – weak acid - which dissociates into an equilibrium solution of hydrated SO2, H+ ions, & Bisulfite (HSO3

-) ionsSO2(g) + H2O(l) [H⇄ 2SO3(aq)] H⇄ +(aq) + HSO3

-(aq)● Neither H2SO3 or H2CO3 (both weak acids) can exist

as isolated molecules – they dissociate immediately in water

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Group 6 – Oxygen Family (ns2np4)● The S in the Sulfite ion (SO3

2-) is in the 4+ state and can be easily oxidized to 6+ state.

● Thus, Sulfites are good Reducing Agents● SO3 (Sulfur Trioxide) is produced by oxidizing SO2

SO2(g) + 1/2O2 SO⇄ 3(g)

● Sulfuric Acid (H2SO4) is a strong acid and the most common industrial chemical It is prepared from SO3, H2O, & conc H2SO4

SO3(g) + conc H2SO4 + H2O H2SO4(l)

● Like other strong acids, sulfuric acid dissociates completely in water forming the Bisulate (HS6+O4)- ion

● Conc Sulfuric acid is an excellent dehydrating agent● The loosely held proton transfer to water in an

exothermic formation of Hydronium ion (H3O+)

V2O5/K2O

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Group 7 – Halogen Family (ns2np5) Trends in properties down the group are just the opposite

of those in Group 1A For Halogens, boiling point, melting point, heats of fusion

& vaporization increase down the group The reason for the opposite trends is the different type of

bonding: Alkali metals exhibit metallic bonding, which decreases

in strength as atoms become larger down group Halogens exist as diatomic molecules that interact

through “Dispersion” forces Halogens are quite reactive reacting with metals and

nonmetals to form ionic and covalent compounds – metal & nonmetal halide oxides, and oxoacids

The halides must gain a single electron to attain the noble gas configuration as a negatively charged anion26

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Group 7 – Halogen Family (ns2np5)

atomic properties

physical properties

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Group 7 – Halogen Family (ns2np5)

Halogen redox behavior is based on electron affinity, ionic charge density, and electronegativity

Halogen higher in group can oxidize halide ion lower in group

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Group 7 – Halogen Family (ns2np5) The reactivity of Halogens decreases down group because

of the decrease in electronegativity

Note: Fluorine is the most electronegative element

The F-F bond is the weakest, despite it short length, because the lone pairs of electrons around the first Fluorine atom repel those on the other Fluorine atom weakening the bond

Because of the weak bond, F2 reacts with every element, except the noble gases, in many cases explosively

The Halogens display the largest range of electronegativity, but all are electronegative enough to behave as nonmetals

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Group 7 – Halogen Family (ns2np5) Halogens act as oxidizing agents (they are reduced, gaining

electrons) in the majority of their reactions Halogens higher in group oxidize Halide ions lower in the

group – Oxidizing ability of X decreases down groupF2(g) + 2X-(aq) 2F-(aq) + X2 (X = Cl, Br, I)

Reaction of Chlorine with water produces Hypochlorous acidCl2 + 2H2O(l) Cl⇄ -(aq) + HClO(aq) + H+ + H2O(l)

Chlorination of drinking water (disinfectant) Household bleach is a dilute 5.25% solution of

Sodium Hypochlorite (NaClO) Hydrogen Chloride (HCL) is extremely water soluble forming

H+ & Cl- ions in solution (Hydrochloric Acid) Hydrochloric Acid occurs in animal stomach fluids and has

many industrial uses

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Group 7 – Halogen Family (ns2np5) Highlights of Halogen Chemistry

The Hydrogen Halides (HX) are formed from the reaction of metal halides and a concentrated acid

CaF2(s) + H2SO4(l) CaSO4(s) + 2HF(g)

2NaBr(s) + H3PO4(l) Na3PO4(s) + 3HBr(g)

HCl is formed as a byproduct in the chlorination of Hydrocarbons for plastics production

CH2=CH2(g) + Cl2(g) ClCH2CH2Cl(g) CH2=CHCl(g) + HCl(g)

Ethylene 1,2 Dichloroethane 1-chloroethene

Hydrogen Fluoride, with its short, strong bond forms a weak acid (Hydrofluoric acid) with water

HF(g) + H2O(l) H3O+ + F-

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Group 7 – Halogen Family (ns2np5) The other Hydrogen Halides (Cl ,Br, I) dissociate

completely to form stoichiolmetric amounts ofH3O+ ions – strong acids

HBr(g) + H2O(l) H3O+(aq) + Br-

(aq)

Halogens react Exothermically with one another to form many “Interhalogen” compounds● Diatomic Molecules (ClF, BrCl, IF)

The more electronegative atom has the 1- charge; the other less electronegative atom has 1+ charge

● The XYn interhalogens (n = 3,4,5) form when the larger members of the group (X) use “d” orbitals to expand the valence shellThe central atom in these molecules has the lower electronegativity and positive charge

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Group 7 – Halogen Family (ns2np5) Commercially useful interhalogens include Fluorine

compounds used as powerful fluorinating agents, reacting with metals, nonmetals, oxides, and even wood and asbestos

Sb(s) + ClF3(l) SnF2(s) + ClF(g)

P4(s) + 5ClF3(L) 4PF3(g) + 3ClF(g)

The Fluoro Interhalogens react very actively (explosively) with water yielding Hydrogen Fluoride (HF) and the oxoacid. There is no oxidation-reduction reaction and the central atom of the oxoacid retains the same oxidation state

3H2O(l) + Br5+F5(l) 5HF(g) + HBr5+O3

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Group 7 – Halogen Family (ns2np5) Oddness & Eveness of Oxidation States

● Odd numbered groups exhibit odd-numbered oxidation states (Na+1, P+5, Cl-1)

● Even numbered groups exhibit even-numbered oxidation states (Ca+2, C+4, O-2)

● Reason: Almost all stable molecules have “Paired” electrons either as bonded or lone pairs

● When bonds form or break, two electrons are involved and the oxidation state changes by “2”

Ex. Consider Interhalogens – XY, XY3, XY5, XY7 With Y in the -1 state, the X atoms must be in the +1, +3, +5, +7 state, respectivelyThe X+1 state arises when Y fills its valence shellThe X+7 state arises when X is completely oxidized(all electrons shifted to more electronegative Y atom

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Group 7 – Halogen Family (ns2np5) Odd Numbered Oxidation States

When I2 reacts with F2, Iodine Fluoride (IF) forms

I2 + F2 2I+1F-1

Each of the two shared electrons in I2 are used to filled the valence shell of each Fluorine

In IF3, Iodine uses two more valence electrons to form two more bonds

I+1F- + F2 I+3F3

If only 1 electron changed, then an unstable lone-electron species containing 2 Fluorines would form

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Group 7 – Halogen Family (ns2np5) Even Numbered Oxidation States

An element in an even-numbered group, such as Sulfur in Group 6A(16), shows the same tendency to have paired electrons in its compounds● Elemental Sulfur (Ox No = 0) gains or shares 2 electrons to

complete its shell● The Sulfur atom loses 2 electrons to react with Fluorine to

form:

SF2 forms a bent compound. There are a total of 20 valence electrons: 6 by Sulfur, 7 from each Fluorine = 20 total. Eight are placed around the sulfur. Six are placed around each Fluorine. A Fluorine is placed on two sides of the Sulfur. The two unshared electron pairs take up more space than the shared pairs and so the shared pairs move closer together approximately 105 degrees apart. AX2E2 = Tetrahedral Bent, just like water.

2SF (Ox No = + 2)

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Group 7 – Halogen Family (ns2np5) Molecular shapes of the main types of interhalogen

compounds

ClF

linear, XY

IF7

Pentagonal Bipyramidal

, XY7

900

ClF3

T-shaped

XY3

BrF5Square Pyramidal

XY5

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Group 7 – Halogen Family (ns2np5) Halogen Oxides

Group 7A Halogens form many Oxides that are powerful oxidizing agents (they are reduced by gaining the electrons lost by the oxidized species)

The Oxides form Acids with water Dichlorine Monoxide (Cl2O) & Chlorine Dioxide (ClO2)

are used to bleach paper2NaClO3(s) + SO2(g) + H2SO4(aq) 2ClO2(g) + 2NaHSO4(aq)

chlorine dioxide

ClO2

ClO2 has an unpaired electron andthe Chlorine (Cl) in the unusual+4 oxidation state(4 electrons are shared with the 2 oxygen atoms, leaving 3 unshaired)

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Group 7 – Halogen Family (ns2np5) Halogen Oxoacids and oxoanions

Oxoacids and Oxoanions are formed by reacting the Halogens and their Oxides with Water

Most Oxoacids are stable only in solution There are four Oxoacids & Oxoanions

Acid Hypochlorous, Chlorous Chloric Perchloric

Salt Sodium Hypochlorite, Sodium Chlorite, Sodium Chlorate, Sodium Perchlorate

The known Halogen Oxoacids

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Group 7 – Halogen Family (ns2np5) Electronegativity of the Halogen

Relative strengths of the Halogen Oxoacids depend on two factors● Electronegativity of the Halogen

The more electronegative the halogen, the more electron density it removes from theO-H bond, and the more easily the proton is lost

Among Oxoacids with the oxidation state of the Halogen in each Halogen the same, the Acidity (acid strength) decreases as the Halogen’s Electronegativity (EN) decreases

Electronegativity – Cl > Br > I Acidity – HOClO2 > HOBrO2 > HOIO2

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Group 7 – Halogen Family (ns2np5)● Oxidation state of the Halogen

The oxidation number is a number identical with the valency but with a sign, expressing the nature of the charge of the species in question when formed from the neutral atom

The oxidation number of Chlorine in Chlorine Oxoacids Hydrochloric Acid (HCl) - 1 Hypochlorous acid (HOCl) + 1 Chlorous Acid (HOCLO or HClO2) + 3 Chloric acid (HClO3 or HOCLO2) + 5 Perchloric acid (HClO4 or HOCLO3) + 7

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Group 7 – Halogen Family (ns2np5)● Oxidation State of the Halogen

Among Oxoacids of a given Halogen, such as Chlorine, acid strength decreases as the oxidation state of Halogen decreases

The higher the oxidation state (also stated as the number of attached O atoms) of the Halogen, the more electron density it pulls from the O-H bond

HOCL+7O3 > HOCL+5O2 > HOCl+3O > HOCl+1

Perchloric Chloric Chlorous Hypochlorous Acid Acid Acid Acid

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Group 8 – Noble Gas Family (ns2np6) The Noble gases have completed outer s & p shells Noble gases are generally not reactive – nearly inert Behave more like “Ideal Gases” than any other gases The smallest radii in their period Condense and solidify only at very low temperatures Helium solidifies (with pressure) at -272.2oC (absolute

zero is -273.15oC) and boils only 3 degrees higher A few Noble gas compounds have been prepared

PtF6 + Xe XePtF6

Other Xenon compoundsXe+2F2 Xe+4F4 Xe+6F6 Xe+8O4

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Group 8 – Noble Gas Family (ns2np6)

atomic properties

physical properties

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Practice Problem What trends exist for Zeff (Effective Nuclear Charge)

across a period and down a group

Ans: Zeff, the effective nuclear charge

In multi-electron atoms an electron feels the attraction from the positively charged nucleus (protons) and the repulsion of like-charged electrons

Electron repulsion shields the electron from the nuclear attraction making the electron easier to remove

Shielding reduces the “full nuclear charge” to an effective nuclear charge (Zeff)

Zeff increases across a period and decreases down a group

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Practice Problem How does the Effective Nuclear Charge, Zeff , influence

atomic size, Ionization Energy (IE), and Electronegativity (EN)?

Ans: As you move to the right across a Period:

The Atomic Size decreases

The Ionization Energy increases

The Electronegativity increases

All because of the increased Zeff

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Practice Problem How are covalent and metallic bonding similar?

Ans: Covalent and Metallic bonding involve sharing of electrons between atoms

Covalent bonding includes sharing between a small number of atoms (usually two)

Metallic bonding involves essentially all the atoms in a given sample

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Practice Problem Which of following pairs react to form covalent

compounds, ionic compounds?

a. Be & C b. Sr & O c. Ca & Cl d. P & F

Ans: a. Covalent

b. Ionic

c. Ionic

d. Covalent

(Oxygen is most ionic in group 6)

(2 non-metals)

Metal & non-metal)

(2 non-metals)

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Practice Problem Which member of each pair gives the more acidic solution? a. CO2 or SrO b. SnO or SnO2

c. Cl2O or Na2O d. SO2 or MgOAns:a. Carbon dioxide will form a more acidic solution 2 CO2(g) + H2O(l) → H2CO3(aq) (Weak acid) SrO(s) + H2O(l) → Sr(OH)2(aq) (Sr oxides form basic solutions) b. Tin(IV) oxide (SnO) An element with more than 1 oxidation

state exhibits less metallic behavior in its higher state – more acidic

c. Dichlorine Oxide (Cl2O) Non-metal oxides form acidic solution; metallic oxides form basic solution

d. Sulfur Dioxide (SO2) Non-metal oxides for acidic solution; metallic oxides form basic solution

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Practice Problem Each of the following properties shows a regular trend in

Group 1A. Predict whether each increases or decreases up the group

a. Melting Point b. E – E bond length c. Hardness

d. Molar Volume e. Lattice Energy of E-Br

Ans: a. Increases

b. Decreases

c. Increases

d. Decreases

e. Increases

Increased Lattice Energy

Decreasing radius of atom

Increased Lattice Energy

Decreasing radius of atom

The atomic radius decreasesas you move up a groupincreasing the lattice energy.

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Practice Problem The melting points of Alkaline Earth metals (group 2A) are

many times higher than the Alkali metals (group 1A)

Explain this difference on the basis of atomic properties

Ans: Metal atoms are held together by metallic bonding, a sharing of valence electrons.

Alkaline earth metal atoms have one more valence electron than alkali metal atoms, so the number of electrons shared is greater

Thus, metallic bonds in alkaline earth metals are stronger than in alkali metals.

Melting requires overcoming the metallic bonds

To overcome the stronger alkaline earth metal bonds requires more energy (higher temperature) than to overcome the alkali earth metal bonds.

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Practice Problem Many compounds of group 3A elements have chemical

behavior that reflects an electron deficiencyExplain electron deficiency with illustrative reactions

Ans: Compounds of Group 3A(13) elements (ns2np1), like Boron (B), have only six electrons in their valence shell when combined with Halogens to form three bonds

Having six electrons, rather than an octet, results in an “electron deficiency,” i.e., violates octet ruleAs an electron deficient central atom, Born (B) is

trigonal planar (AX3). Upon accepting an electron pair to form a bond, the shape changes to tetrahedral (AX4)

BF3(g) + NH3(g) → F3B–NH3(g)

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Practice ProblemNearly every compound of Silicon (Group 4A) has the element in the +4 oxidation state. In contrast, most compounds of Lead have the element in the +2 state a. What general observation does this fact illustrate? Ans: The increased stability of the lower oxidation state as one goes down a group b. Explain in terms of atomic structure and molecular properties Ans: As the atoms become larger (Pb > Si), the strength of the bonds to other elements becomes weaker, and insufficient energy is gained in forming the bonds to offset the additional ionization or promotion energy c. Give an analogous example from Group 3A Ans: Thallium(Tl+) is more stable than Tl3+, but Al3+ is the only table oxidation state for Al

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Practice Problem

Based on the relative sizes of Fluorine (F) and Chlorine(Cl), predict the structure of PF2Cl3

Ans: From the Lewis structure, the Phosphorus (Central atom) has 5 electron groups for a trigonal bipyramidal molecular shape

In this shape, the three groups in the equatorial plane have greater bond angles (120°) than the two groups above and below this plane (90°)

The Chlorine atoms (larger than Fluorine atoms) would occupy the planar sites where there is more space for the larger atoms

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Practice Probem A Halogen (X2) disproportionates (one X is reduced and one is

oxidized) in base in several steps to X- and XO-3. Write the overall reaction for the disproportionation of Br2 to Br- and BrO3

-1

Ans: A substance that disproportionates serves as both an oxidizing and reducing agent.

Assume that OH– serves as the base

Write the reactants and products of the reaction, and balance like a redox reaction

Br2(l) + OH– (aq) → Br– (aq) + BrO3-1 (aq) + H2O(l)

Br2(l) + 6 OH– (aq) → Br– (aq) + BrO3-1 (aq) + 3 H2O(l)

3 Br2(l) + 6 OH– (aq) → 5 Br– (aq) + BrO3-1 (aq) + 3 H2O(l)

5 x -1 = -5 +50

Balance e- on each side using coefficients

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Practice Problem The main reason Alkali metal dihydrides (MX2) do not

form is the Ionization Energy (IE) of the metal

Why is the IE so high for alkali metals

Ans: Alkali metals have an outer electron configuration of ns1

The first electron lost by the metal is the ns electron, giving the metal a noble gas configuration

Second ionization energies for alkali metals are high because the electron being removed is from the next lower energy level and electrons in a lower level are more tightly held by the nucleus.

The metal would also lose its noble gas configuration