Bonding and Periodic Table

TrendsHonors Chemistry

Electron Configurations

• Stable Octet: 8 electrons8 electrons in the outer level is very stable (includes He)

• Ions – gain/lose electrons to achieve a stable octet

• Isoelectronic – same electron configuration• Examples: N, O, F, Na, Mg, Al are isoelectronic

with Ne – this is called an isoelectronic series

• Pseudoisoelectronic – same electron configuration but includes the d orbitals

• Fe+2 is pseudoisoelectronic with Ar

Periodic Table TrendsIntroduction• Properties are periodic• An element’s position & its properties are a

result of its electrons• The outermost electrons, aka valence

electrons, have the greatest influence on the properties of the elements.

• Adding an electron to an inner core orbital results in less striking changes in properties than adding an electron to an outer valence orbital (higher energy).

• Shielding Effect: electrons in the lower energy levels (inner core electrons), shield electrons in the outer levels from the full effect of the nuclear charge.

Trends in the Periodic TableA. Atomic Radius

1. The distance from the center of the nucleus to the outermost electron.

2. Bond Radius3. Atoms get larger

going down a group and smaller going across a period.

Ex) Na is larger than Mg Na is smaller than K Ga vs. Al

Atomic Radii of the Representative Elements

Atomic Radii vs Atomic Number

Positive Ion Size

• When atoms lose electrons, they become (positive) and get smaller.

• The sizes of cations increases down a group.

• The sizes of cations decreases across a period.

Negative Ionic Size

• When atoms gain electrons, they become (negative) and get larger.

• The sizes of anions increases down a group.

• The sizes of anions decreases across a period.

Relative Sizes of Positive &

Negative Ions

The sodium ion lost an electron, and therefore the positive

protons in the nucleus exert a stronger pull on the remaining negative electrons, shrinking

the orbitals. Thus positive ions are smaller than their atoms.

The chloride ion gained an electron, and therefore the fewer positive protons in the nucleus

exert a weaker pull on the extra negative electrons, increasing the size of the orbitals. Thus negative ions are larger than

their atoms.

• Within an isoelectronic series, radii decrease with increasing atomic number because of increasing nuclear charge.

N-3 O-2 F-1 Na+1 Mg+2 Al+3

How many electrons? Nuclear charge?

> > > > >

Ionic Radius

• Cations & anions decrease in size going across a period

• Cations & anions increase in size going down a group

Electron Attraction in a Bond &

Ion Size

Ionization Energy (IE):

1. The energy needed to remove one electron from an atom. (kJ/mole)

2. IE measures how tightly electrons are bound to an atom.

– Elements that do not want to lose their electrons have high ionization energies.

– Elements that easily lose electrons have low ionization energies.

X + energy X + energy X X+1+1 + 1 e + 1 e--

1st Ionization Energy

3. I.E. decreases down a group. 4. I.E. increases across a period

– Account for deviations across a period.

5. Metals tend to have low IE1.

6. Nonmetals tend to have high IE1.

Ionization Energy of the 1st 20 Elements

1. Energy required to remove electrons beyond the 1st electron.

2. Ionization energies will increase for every electron removed.

X + IEX + IE11 X X+1+1 + 1 e + 1 e- -

XX+1+1 + IE + IE22 X X+2+2 + 1 e + 1 e--

XX+2+2 + IE + IE33 X X+3+3 + 1 e + 1 e--

Successive Ionization Energies:

Successive Ionization Energies:

kJ/mol IE1 IE2 IE3 IE4 IE5

Na 496 4562 6912 954413

353

Mg 738 1451 773310 540

13 628

Al 578 1817 274511 578

14 831

3s

3s

Electron Configuration

Na: [Ne] 3s1

Mg: [Ne] 3s2

Al: [Ne] 3s23p13p3s

Ionization Energy vs. Atomic Number

Notice the dips across the period… why?

Notice the dips across the period… why?

Period 3

Na - [Ne] 3s1 ___ ___ ___ ___Mg - [Ne] 3s2 ___ ___ ___ ___Al - [Ne] 3s23p11 ___ ___ ___ ___Si - [Ne] 3s23p22 ___ ___ ___ ___P - [Ne] 3s23p33 ___ ___ ___ ___S - [Ne] 3s23p44 ___ ___ ___ ___Cl - [Ne] 3s23p55 ___ ___ ___ ___Ar - [Ne] 3s23p66 ___ ___ ___ ___

3s 3p

Electronegativity (EN)

1. Reflects an atoms ability to attract electrons in a chemical bond.• Up to 4.0 for F• Zero for He, Ne, Ar and Kr

2. Metals have low EN.3. Nonmetals have high EN.4. EN decreases down a group.5. EN increases across a period.

Electron Affinity (EA)1. Energy change that occurs when a neutral

gaseous atom gains an electron. Units kJ/mol.1 e- + X X-1 + EA

• Most elements have no affinity for an additional electron and have an EA equal to zero.

He(g) + e- He- EA = 0 kJ/mol

He will not add an electronCl(g) + e- Cl- + 349 kJ/mol EA = -349 kJ/mol

Exothermic!!!

Electron Affinity (EA)

2. Metals have low EA.3. Nonmetals have high EA.4. EA decreases down a group.5. EA increases (becomes more

negative) across a period.• EXCLUDES noble gases• Exceptions: Groups IIA (~0) and VA

(~0 for N and smallerfor P to Bi)• Why? Filled s and half filled p

Metallic Character

1. Reflected by those elements that can lose electrons easily.

2. Increases down a group.3. Decreases across a period.4. The most metallic metal is Cesium.5. The most nonmetallic (least

metallic) metal is Aluminum.

Reactivity

• Related to the ability of an element to lose or gain an electron

Brainiac Alkali Metals Alkali Metals Reactivity

1. Reactivity of metals INCREASE down a group.

2. Reactivity of metals DECREASE across a period.

3. Reactivity of nonmetals DECREASE down a group.

4. Reactivity of nonmetals INCREASE across a period.

Chemical BondingChapter 6

Honors Chemistry

8

Introduction

• A chemical bond is an attractive force that holds atoms together in elements or compounds (to function as a unit)– intramolecular (within) vs. intermolecular (between)

• Bond energy is the energy needed to break or form 1 mole of bonds in a gaseous substance (kJ/mol)

• Bonding usually involves only the valence electrons.– In most compounds of the representative elements, the

atoms have an electron configuration that is isoelectronic or psuedoisoelectronic with a noble gas

• The manner in which atoms are bound together in a given substance has a profound effect on its chemical and physical properties.

Octet Rule

• Many chemical compounds form such that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level.– Very stable

– There are exceptions: H, He, B (BF3)

• System achieves the lowest possible energy.

Types of Chemical Bonds:

Metallic BondsIonic Bonds

Covalent Bonds

Metallic Bonds

• Simplest crystalline solid – arranged in a very compact and orderly pattern

• Sea of electrons – the valence electrons are mobile around metal cations – Electrons are delocalized

• Attraction of the metal atoms and the surrounding sea of electrons

Metallic Bonds

• Explains metallic properties:– High electrical and thermal conductivity

(flow of electrons)– Luster (metals absorb wide range of -

excites e- and fall back emitting E in form of light results in shiny appearance)

– Ductility & malleability (mobility of e-, metallic bonding is same in all directions throughout solid)

Metallic bonding visual on Holt

Metals vs. Ionic Crystals

• Metallic properties due to sea of electrons

• Ionic compounds are hard but brittle – repulsions result from shift and causes crystal to break

• Chemical bonding that results from an electrostatic attraction between cations and anions to form a neutral compound.

• “salts”• Octet Rule

a) Atoms will transfer electrons (e-) to each other in order to have a full set of valence electrons.

b) When electrons are transferred, ionic bonds are formed.

Ionic bonding

Ionic bonding

Covalent Bonding

• Sharing one or more electron pairs between 2 atoms

Characteristics of Ionic Compounds

(hundreds of compounds)1. All are high melting solids (>400°C).

a) Orderly 3D arrangements (pattern) called crystalline solid or crystal lattice.

b) Simplest arrangement = formula unit

c) High mp reflects strong bonds – large attractive forces are very stable

d) Many are white.e) Colored compounds usually

contain the transition elements (Cu, Cr, Co, Ni, Mn)

2. Solubilitya) Many are soluble in polar solvents, such as

water (aka aqueous solutions)b) Most are insoluble in nonpolar solvents,

such as hexane (C6H14)

Characteristics of Ionic Compounds:

Dissolving Salt Animation

Characteristics of Ionic Compounds:

3. Conductivitya) Solids are non

conductive – ions cannot move freely

b) Molten compounds are conductive – ions move freely (NaCl mp ~800°C)

c) Aqueous solutions are conductive – ions free to move in solution

Formation of Ionic Compounds

• Formation results from a transfer of electrons and the electrostatic attractions of the closely packed, oppositely charged ions.

• Ionic substances are formed when an atom that loses electrons relatively easily reacts with an atoms that has a high affinity for electrons.

• Forms between a metal and a nonmetal. (large EN difference)– Metal: low IE, EN, EA– Nonmetal: high IE, EN, EA– Metal is oxidized (loss of e-) and

nonmetal is reduced (gain of e-)• The ion pair has lower energy than

separated ions.

Electron Configuration

Distribution of electron density• Na: 1s22s22p63s1

– 186 pm

• Cl: 1s22s22p63s23p5

– 99 pm

• Na+1: 1s22s22p6

– 95 pm

• Cl-1: 1s22s22p63s23p6

– 181 pm

Lewis structures or electron-dot structures

Lewis structures examples:

• Sodium and chlorine

• Potassium and phosphorus

Ionic, Nonpolar Covalent, Polar Covalent

Character of Bonds

ENDifference 0.00 0.65 0.94 1.19 1.43 1.67 1.91 2.19 2.54 3.03

% IonicCharacter 0% 10% 20% 30% 40% 50% 60% 70% 80% 90%

% CovalentCharacter 100% 90% 80% 70% 60% 50% 40% 30% 20% 10%

HOLT

Nonpolar Covalent Bonds

• Electron pair is shared equally between the atoms (ΔEN = 0 to ~0.4)– Diatomic molecules (H O F Br I N Cl); allotropes

(S8)

Nonpolar Covalent BondsHydrogen

• Electron pair is shared unequally between atoms (ΔEN = ~0.4 to ~1.9) • Results in an electric dipole (2 poles)• Equal but opposite charges that are separated

by a short distance• Separation of charge between 2 covalently

bonded atoms

• Examples: HF, HBr, H2O

Polar Covalent Bonds

Polar and Nonpolar Bonds

1. They are gases, liquids, or solids with low melting points (<300°C)

2. Usually the simplest arrangement is called a molecule (many sizes and shapes)

Characteristics of Covalent Compounds

(~11 million compounds)

3. Solubility– Many are insoluble in polar solvents– Many are soluble in nonpolar solvents

4. Conductivity– Liquid and molten compounds are

nonconductors– Aqueous solutions are usually

nonconductors or poor conductors

Characteristics of Covalent Compounds

Formation

1. Formation due to sharing electrons

2. Often forms between nonmetals

– metalloid and nonmetals

– small electronegativity differences

Potential Energy Changes

a) PE=0 separation of atoms do no affect each other

b) PE decreases as atoms are drawn together by attractive forces

c) PE is at MINIMUM when attractive forces are balanced by repulsive forces = stable

d) PE is increasing when repulsion between like charges outweighs attraction between opposite charges

Potential Energy Changes

HOLT

Bond length & energy

•Bond Length = distance between 2 nuclei in most stable position (lowest PE)– 1s orbitals overlap (head on) to

form a single covalent bond (sigma bond)

•Bond Energy = energy required to break a bond

3. Number of covalent bonds likely to form for nonmetals or metalloids depends upon the number unpaired electrons

4. Electron configurationsCl2: 1s22s22p63s23p23p23p1

1s22s22p63s23p23p23p1

1 bond: share 1 pair of e- = single bond

Formation

Electron config for covalent compounds

O2 1s22s22p22p12p1

1s22s22p22p12p1

2 bonds: share 2 pairs of e- = double bond

7 Elements that can multiple bond: C, N, O, Si, P, S, Se

Bond Energy for Carbon Bonds

• Bond lengths (Å):Single > double > triple 1.54 1.34 1.20

• Bond Energy (kJ/mol) Single < double < triple 346 612 835

Multiple Bonds

Lewis StructuresGuidelines:1. Select a reasonable skeleton for the molecule

or polyatomic ion. The central atom is the least electronegative atom (excluding H)

2. Calculate the number of shared electrons (S):

S = N – A– N = total number of valence electrons required (all

8, except H)– A = number of valence electrons available

3. Place shared pairs of electrons in skeleton4. Place lone pairs (for octets)5. NOTE: # of “dots” in the Lewis structure = A

Examples of Lewis Structures for Covalent

Bonds• Cl2

• CO2

• ClO4-1

Structural formula (line structure) only shows how the molecule or polyatomic ion is bonded – NO “dots” shown

Coordinate Covalent Bonds (Dative bonds)

• A bond formed when 1 atom provides both electrons (to covalent bond)

• Perchlorate (ClO4-1) has 4 coordinate

covalent bonds

Equivalent Lewis Structures

• Also known as resonance structures

• CO3-2

Limitations of the Octet Rule

1. Less than an octet (electron deficient)

• Be (4) & B (6)

• BF3

2. More than an octet (expanded valence shell or hypervalence)

• 8 elements: P, S, As, Br, Sb, Te, I, Xe

• SF4

3. Odd number of electrons• NO

Overall, if you use N-A-S…

• If S leads to too many bonds– 1st - look at central atom. Is it electron

deficient? (B, Be)– 2nd - multiple bond

• If S lead to too few bonds– Think hypervalence

Molecular Shapes

1. Valence shell electron pair repulsion theory (VSEPR Theory): helps to predict the spatial arrangement of atoms in a molecule or polyatomic ion.

a) Introductioni. The central atom is any atom bonded to more

than one other atom.ii. Unshared pairs (lone pair) of electrons and

bonding pairs on the central atom orient themselves to minimize repulsions.

iii. Lone pairs of electrons occupy MORE space than bonding pairs.

VSPER

b. Counting regions of high electron density around the central atom.

i. Each bonded atom is counted as ONE region of high electron density, whether it is a single, a double, or a triple bond.

ii. Each unshared pair of valence electrons on the central atom is counted as ONE region of high electron density.

Valence Bond (VB) Theory

• Describes HOW bonding occurs• Usually atomic orbitals do not have

the correct energies or orientation to describe where the electrons are when bonded to other atoms

• Hybridization is the mixing of the atomic orbitals to form new hybrid orbitals (s – p – d)

HOLT VSPER

Regions of High Electron Density

• Two regions – LINEARLINEAR arrangement• 2 regions e- density = sp• Bond angle = 180°

• Example: BeH2

Animation of sp hybridization

Regions of High Electron Density

• Three regions – TRIGONAL PLANARTRIGONAL PLANAR arrangement

• 3 regions e- density (all bonding) = sp2

• Bond angle = 120°

• Example: BF3

Animation of sp2 hybridization

Regions of High Electron Density

• 3 regions e- density = sp2 – 2 bonding and 1 lone pair

• Electronic geometry – trigonal planar• Molecular geometry – BENTBENT or

ANGULARANGULAR• Bond angle = 115°• Example: NOCl

Regions of High Electron Density

Regions of High Electron Density

• Four regions – TETRAHEDRALTETRAHEDRAL arrangement

• 4 regions e- density (all bonding) = sp3

• Bond angle = 109.5°

• Example: CH4

Regions of High Electron Density

sp3 hybrid orbital

Animation of sp hybridization

Regions of High Electron Density

• 4 regions e- density = sp3

– 3 bonding and 1 lone pairs

• Electronic geometry – tetrahedral• Molecular geometry – TRIGONAL TRIGONAL

PYRAMIDALPYRAMIDAL • Bond angle = 107.3°

• Example: NH3

Regions of High Electron Density

Regions of High Electron Density

• 4 regions e- density = sp3

– 2 bonding and 2 lone pairs

• Electronic geometry – tetrahedral• Molecular geometry – BENTBENT or

ANGULARANGULAR• Bond angle = 104.5°

• Example: H2O

Regions of High Electron Density

Molecular Polarity

Consider…a) The presence of at least 1 polar

bond or 1 lone pair of electrons andb) The molecular shape to determine

the overall molecular polarityExamples:HCl, BeCl2, BF3, CH4, NH3, H2O

Molecular Polarity

http://preparatorychemistry.com/Bishop_molecular_polarity.htm

Intermolecular Forces(aka Van der Waals forces)

• Forces of attractions BETWEEN molecules– Intermolecular forces are weaker than

intramolecular forces

• Important for states of matter• Boiling point is a measure of the

strength of these– High bp – strong intermolecular forces

Intermolecular Forces(aka Van der Waals forces)

1. London Dispersion Forces • Attractions caused by temporary

(instantaneous) dipoles and are present between ALL atoms and molecules

• Increases with size• Holt

He Ne Ar Kr

Boiling Point

-269° -246° -186° -153°

London Dispersion Forces: induced dipole–induced dipole interactions

Intermolecular Forces(aka Van der Waals forces)

2. Dipole-Dipole Forces • Attractions between polar molecules• Permanent dipoles• Holt

F2

(nonpolar)

HCl(polar)

BrF(polar)

CF4

(nonpolar)

CH3F(polar)

Boiling Point -188° -89° -20° -128° -84 °

3. Hydrogen Bonding– Attractions resulting when

hydrogen that is bonded to a highly EN atom (F, N, O) is attracted to an unshared pair of electrons of an EN atom in a nearby molecule

– H2S vs H2O

bp -59.6° 100°– NH3 vs CH4

bp -33° -161°

Intermolecular Forces(aka Van der Waals forces)

For each of the molecules below, list the types of intermolecular forces which act

between pairs of these molecules.

(a) CH4

(b) PF3

(c) CO2

(d) HCN(e) HCOOH (methanoic acid)

(a)CH4 is a tetrahedral molecule - it does not have a permanent dipole moment. – The figure above shown CH4 in two views: one shows it as

it is commonly drawn, with one H at the top and three H's at the bottom. The second figure shows CH4 rotated to fit inside a cube. This might help to make clear why it does not have a permanent dipole moment. The dipole moments of the two C-H bonds pointing up exactly cancel the dipole moments of the two C-H bonds pointing downward.

• CH4 does not contain N, O, or F and therefore there are no hydrogen bonds between CH4 molecules.

• Therefore only dispersion forces act between pairs of CH4 molecules.

• Other tetrahedral molecules like CF4 and CCl4 do not have a permanent dipole moment.

website

Next: PF3

(b) PF3 is a trigonal pyramidal molecule (like ammonia, the P has a single lone pair of electrons); it does have a permanent dipole moment.

• It does contain F, but it does not contain any hydrogen atoms so there is no possibility of forming hydrogen bonds.

• Therefore dispersion forces and dipole-dipole forces act between pairs of PF3 molecules.

Next: CO2

(c) CO2 is a linear molecule; it does NOT have a permanent dipole moment

• It does contain O, however the oxygen is not bonded to a hydrogen.

• Therefore only dispersion forces act between pairs of CO2 molecules.

Next: HCN

(d) HCN is a linear molecule and it does have a permanent dipole moment

• It does contain N, however the nitrogen is not directly bonded to a hydrogen.

• Therefore dispersion forces and dipole-dipole forces act between pairs of HCN molecules.

Next: HCOOH

e) HCOOH is a non-linear molecule; it does have a permanent dipole moment;

• It does contain O, and the oxygen is directly bonded to a hydrogen.

• Therefore dispersion forces, dipole-dipole forces and hydrogen bonds act between pairs of HCOOH molecules.

The intermolecular forces acting between pairs of these molecules.

(a) CH4

(b) PF3

(c) CO2

(d) HCN(e) HCOOH

(methanoic acid)

(a) dispersion(b) dispersion, dipole-

dipole(c) dispersion(d) dispersion, dipole-

dipole(e) dispersion, dipole-

dipole, hydrogen bonds