Periodic Tableand

Atomic Theory

Elements and Chemical Symbols

• An element is a type of matter that cannot be broken down into two or more pure substances.– Some facts about elements:

• There are 112 known elements

• 91 occur naturally

• Many are familiar to us:– Charcoal (for bbq’s) is nearly pure carbon

– Aluminum is used on many household utensils

– Thermometers use mercury (liquid metal)

The Periodic Table

• The period table was developed by Mendeleev.

– by arranging elements according to atomic mass, he discovered the Periodic Law which is the foundation of the Periodic Table

• Periodic Law - patterns of reactivity, behavior, and properties occur in families or groups of elements.

Patterns in the Periodic Table

• Elements are organized according to four basic patterns:

1. atomic number

2. metals & non-metals

3. period

4. group/family

1. Atomic Number

• refers to the number of protons in an atom. Protons are positively charged subatomic particles in the atom’s nucleus.

2. Metals & Non-metals

• Metals: • left of staircase

• good conductors of heat and electricity

• shiny

• malleable & ductile

• form positive ions

• all except mercury are solids at SATP

• Non-metals:• right of staircase

• poor conductors

• non shiny

• brittle (solids)

• form negative ions

• either solids, liquids, or gas at SATP

3. Periods

• are the horizontal rows of elements whose properties change from metallic to non-metallic from left to right along the row.

4. Groups/Families

• are the vertical columns in the main part of the table whose elements have similar chemical properties. They are numbered left to right #1 - 18.

Group 1: alkali metals

– react violently with water and get more reactive as you move DOWN the group

– hydrogen is not a member.

Group 2: alkaline earth metals

– light, reactive metals

Group 17: halogens

– extremely reactive non-metals

– F2(g) is most reactive

Group 18: noble gases

– extremely low chemical reactivity or inert (very stable)

Metalloids

– Have properties of both metals and non-metals. They lie along the staircase of the periodic table.

• Lanthanides: atomic numbers 57-71

• Actinides: atomic numbers 89-103

• Why do elements in the same group have similar properties?

• What is it about metals and non-metals that allow you to predict the compounds they will form?

»ELECTRONS

• Valence Electrons: electrons in the outer energy level of an atom. Valence electrons determine the chemistry of that atom; knowing the number of valence (outer) shell electrons help predict the formation of compounds, name the compounds, and write their chemical formulas.

• Octet rule - we consider an atom with eight electrons in the outer level to have a full outer level. One of the basic rules in chemistry is that an atom with eight electrons in its outer level is particularly stable (H and He are exceptions).

Atomic Theory

History of the atom

•Democritus (400 BC) suggested that the material world was made up of tiny, indivisible particles

• atomos, Greek for “uncut”

• Aristotle believed that all matter was made up of 4 elements, combined in different proportions

• Fire - Hot

• Earth - Cool, heavy

• Water - Wet

• Air - Light

• The “atomic” view of matter faded for centuries, until early scientists attempted to explain the properties of gases

Re-emergence of Atomic Theory

John Dalton postulated that:

1. All matter is composed of extremely small, indivisible particles called atoms.

2. All atoms of a given element are identical (same properties); the atoms of different elements are different.

3. Compounds are formed when atoms of more than one element combine

• A given compound always has the same relative number and kind of atoms

4. Atoms are neither created nor destroyed in chemical reactions, only rearranged.

Make-up of the Atom

• By the 1850s, scientists began to realize that the atom was made up of subatomic particles

• Thought to be positive and negative

Cathode Rays and Electrons

• Mid-1800’s scientists began to study electrical discharge through cathode-ray tubes. Ex: neon signs

• Partially evacuated tube in which a current passes through

• Forms a beam of electrons which move from cathode to anode

• Electrons themselves can’t be seen, but certain materials fluoresce (give off light) when energised

• JJ Thompson observed that when a magnetic or electric field are placed near the electron beam, they influence the direction of flow

• Electrons flow from the negative electrode (cathode) to positive electrode (anode), so electrons are (-)ve.

• http://www.chem.uiuc.edu/clcwebsite/video/Cath.mov

• Magnetic field forces the beam to bend depending on orientation

• Thompson concluded that:

• Cathode rays consist of beams of particles

• The particles have a negative charge

• Thompson understood that all matter was inherently neutral, so there must be a counter

• A positively charged particle, but where to put it

• It was suggested that the negative charges were balanced by a positive umbrella-charge

• “Plum pudding model” “chocolate chip cookie model”

Rutherford and the Nucleus

• This theory was replaced with another, more modern one

• Ernest Rutherford (1910) studied angles at which alpha (a) particles were scattered as they passed through a thin gold foil

Rutherford expected …

• Rutherford believed that the mass and positive charge was evenly distributed throughout the atom, allowing the aparticles to pass through unhindered

a particles

Rutherford explained …

+

• Atom is mostly empty space

• Small, dense, and positive at the center

• Alpha particles were deflected if they got close enough

a particles

• Nucleus: Containing protons and neutrons, it is the bulk of the atom and has a positive charge associated with it

• Electron cloud: Responsible for the majority of the volume of the atom, it is here that the electrons can be found orbiting the nucleus

The modern atom is composed of two regions:

Niels Bohr’s Model of the Atom

• Fixed energy related to the orbit

• Electrons cannot exist between orbits

Bohr’s Model

Nucleus

Electron

Orbit

Energy Levels

Nucleus

Electron

Orbit

Energy Levels

The Bohr Model of the Atom

• Why don’t the electrons fall into the nucleus?

• Move like planets around the sun.

• In circular orbits at different levels.

• Amounts of energy separate one level from another.

How did he develop his theory?

• He used mathematics to explain the visible spectrum of hydrogen gas

Radio

waves

Micro

waves

Infrared

.

Ultra-

violet

X-

Rays

Gamma

Rays

Low

energy

High

energy

Low

FrequencyHigh

FrequencyLong

WavelengthShort

WavelengthVisible Light

The line spectrum

• electricity passed through a gaseous element emits light at a certain wavelength

• Can be seen when passed through a prism

• Every gas has a unique pattern (color)

Line spectrum of various elements

• When an atom or a molecule absorbs some electrical energy, one (or more) of its electrons gains energy and goes to a higher energy level.

• Electrons in this excited state return to the lower energy level and release this excess energy in the form of light.

• The diffraction grating breaks up the predominant colour into bright lines of specific colours that represent the electron transitions occurring.

• The different sized energy jumps are represented by the different colours.

• Bohr’s theory did not explain or show the shape or the path traveled by the electrons.

• His theory could only explain hydrogen and not the more complex atoms

Drawback to Bohr’s Model

Bohr-Rutherford Electron Diagrams

• The purpose of these diagrams is to show the nucleus with its correct # of protons and neutrons, but more importantly, to give a more detailed view of how the electrons are arranged.

1. Draw the nucleus as a solid circle.

2. Put the number of p+ in the nucleus with the number of n0 under it.

3. Place the number of e- in orbits around the nucleus by drawing circles around the nucleus.

• When placing electrons in energy levels you have to fill the innermost shell before beginning to fill the next energy level.

Bohr Rutherford Diagram

18p

36n

Nucleus with

protons and

neutrons

Electron Orbits

(shells) with a

2,8,8 pattern

22n

• The maximum number of electrons in the first energy level is 2.

• The maximum number of electrons in the second energy level is 8.

Ex. Hydrogen would have 1 electron in the first level.

• Valence Electrons are the electrons found in the outermost shell of the atom

• Valence electrons are important in determining how an element reacts chemically with other elements.

• An element with a full valence shell is stable

Practice

• Draw all the Bohr-Rutherford diagrams for the odd numbered elements from 1-19

• Worksheet

– Complete table on Isotope page

– Complete “Which Atom is Which’

Subatomic Particles & Isotopes

Major Subatomic Particles

• Atoms are measured in picometers, 10-12 meters• Hydrogen atom, 32 pm radius

• Nucleus tiny compared to atom• If the atom were a stadium, the nucleus would be a marble

• Radius of the nucleus is on the order of 10-15 m

• Density within the atom is near 1014 g/cm3

Name Symbol Charge Relative Mass (amu)

Actual Mass (g)

Electron e- -1 1/1840 9.11x10-28

Proton p+ +1 1 1.67x10-24

Neutron no 0 1 1.67x10-24

Elemental Classification

• Atomic Number (Z) = number of protons (p+) in the nucleus

• Determines the type of atom

• Mass Number (A) = number of protons + neutrons [Sum of p+ and nº]

• In a neutral atom there is the same number of electrons (e-) and protons (atomic number)

Nuclear Symbols

EA

Z

elemental symbol

mass number

atomic number

• Find the

• number of protons

• number of neutrons

• number of electrons

• atomic number

• mass number

F19

9Br

80

35

Isotopes

• Atoms of the same element can have different numbers of neutrons and therefore have different mass numbers

• The atoms of the same element that differ in the number of neutrons are called isotopes of that element

Hydrogen – 1 Hydrogen - 2 Hydrogen - 3

• When naming, write the mass number after the name of the element

H1

1Protium

H2

1Deuterium

H3

1Tritium

• We are more concerned with average atomic masses, rather than exact ones

• Based on abundance of each isotope found in nature

• We can’t use grams as the unit of measure because the numbers would be too small

• Instead we use Atomic Mass Units (amu)• Standard amu is 1/12 the mass of a carbon-12 atom

• Each isotope has its own atomic mass

Ions

• Ions are formed when an atom gains or loses electrons to become stable.

• Ex . Potassium loses one electron and becomes the ion K+.

Sulfur gains two

S S-2

Atom Ion

• If not told otherwise, the mass of the isotope is the mass number in amu

• The average atomic masses are not whole numbers because they are an average mass value

• Remember, the average atomic masses are the decimal numbers on the periodic table

Average Atomic Masses

• If an element has an atomic number of 34 and a mass number of 78 what is the:

• number of protons in the atom?

• number of neutrons in the atom?

• number of electrons in the atom?

• complete symbol of the atom?

• If an element has 91 protons and 140 neutrons what is the:• atomic number?

• mass number?

• number of electrons?

• complete symbol?

• Alkali metals all end in s1

• Alkaline earth metals all end in s2

• really have to include He but it fits better later.

• He has the properties of the noble gases.

s2

s1 S- block

Transition Metals -d block

d1d2 d3

s1

d5 d5 d6 d7 d8s1

d10 d10

The P-block p1 p2 p3 p4 p5 p6

F - block

• inner transition elements

f1 f5f2 f3 f4 f6 f7 f8 f9 f10 f11 f12 f14f13

Atomic Size

Atomic Size

• The electron cloud doesn’t have a definite edge.

• They get around this by measuring more than 1 atom at a time.

• Summary: it is the volume that an atom takes up

Atomic Size

•Atomic Radius = half the distance between two nuclei of a diatomic molecule.

}Radius

Group trends

• As we go down a group (each atom has another energy level) the atoms get bigger, because more protons and neutrons in the nucleus

H

Li

Na

K

Rb

Periodic Trendsatomic radius decreases as you go from left to right across a

period.

• Why? Stronger attractive forces in atoms (as you go from left to right) between the opposite charges in the nucleus and electron cloud cause the atom to be 'sucked' together a little tighter.

Na Mg Al Si P S Cl Ar

Reactivity

Reactivity

• Reactivity refers to how likely or vigorously an atom is to react with other substances. This is usually determined by how easily electrons can be removed (ionization energy) and how badly they want to take other atom's electrons

For Metals:

Period - reactivity decreases as you go from left to right across a period. Group - reactivity increases as you go down a group

• Why? The farther to the left and down the periodic chart you go, the easier it is for electrons to be given or taken away, resulting in higher reactivity

For Non-metals

• Period - reactivity increases as you go from the left to the right across a period. Group - reactivity decreases as you go down the group.

• Why? The farther right and up you go on the periodic table, the higher the electronegativity, resulting in a more vigorous exchange of electron.

Ionization Energy

Ionization Energy

• The amount of energy required to completely remove an electron from a gaseous atom.

• An atom's 'desire' to grab another atom's electrons.

• Removing one electron makes a +1 ion.

• The energy required is called the first ionization energy.

X(g) + energy →X+ + e-

Ionization Energy

The second and third ionization energies can be represented as follows:

• X+(g) + energy X2+

(g) + e-

• X2+(g) + energy X3+

(g) + e-

• More energy required to remove 2nd electron, and still more energy required to remove 3rd

electron

Period trend

Electronegativity increases as you go from left to right across a period.

• Why? Elements on the left of the period table have 1 -2 valence electrons and would rather give those few valence electrons away (to achieve the octet in a lower energy level) than grab another atom's electrons. As a result, they have low electronegativity. Elements on the right side of the period table only need a few electrons to complete the octet, so they have strong desire to grab another atom's electrons.

Group Trend

electronegativity decreases as you go down a group.

• Why? Elements near the top of the period table have few electrons to begin with; every electron is a big deal. They have a stronger desire to acquire more electrons. Elements near the bottom of the chart have so many electrons that loosing or acquiring an electron is not as big a deal.

• This is due to the shielding affect where electrons in lower energy levels shield the positive charge of the nucleus from outer electrons resulting in those outer electrons not being as tightly bound to the atom.

Shielding

• Shielded slightly from the pull of the nucleus by the electrons that are in the closer orbitals.

Group trends

• Ionization energy decreases down the group.

Going from Mg to Be, IE decreases because:

– Be outer electron is in the 3s sub-shell rather than the 2s. This is higher in energy

– The 3s electron is further from the nucleus and shielded by the inner electrons

– So the 3s electron is more easily removed

• A similar decrease occurs in every group in the periodic table.

Period trendsIE generally increases from left to right.

• Why?

• From Na to Ar (11 protons to 18 protons), the nuclear charge in each element increases.

• The electrons are attracted more strongly to the nucleus – so it takes more energy to remove one from the atom.

Why is there a fall from Mg to Al?

• Al has configuration 1s2 2s2 2p6 3s2 3p1, its outer electron is in a p sublevel

• Mg has electronic configuration 1s2 2s2 2p6 3s2.

• The p level is higher in energy and with Mg the s sub level is full – this gives it a slight stability advantage

Why is there a fall from P to S?

• This can be explained in terms of electron pairing.• As the p sublevel fills up, electrons fill up the vacant

sub levels and are unpaired.• This configuration is more energetically stable than S

as all the electrons are unpaired. It requires more energy to pair up the electrons in S so it has a lower Ionisation energy.

• There is some repulsion between the paired electrons which lessens their attraction to the nucleus.

• It becomes easier to remove!

Driving Force

• Full Energy Levels are very low energy.

• Noble Gases have full energy levels.

• Atoms behave in ways to achieve noble gas configuration.

2nd Ionization Energy

• For elements that reach a filled or half filled sublevel by removing 2 electrons 2nd IE is lower than expected.

• Makes it easier to achieve a full outer shell

• True for s2

• Alkaline earth metals form +2 ions.

Electronegativity

Electronegativity

• The tendency for an atom to attract electrons to itself when it is chemically combined with another element.

• How fair it shares.

• Big electronegativity means it pulls the electron toward it.

• Atoms with large negative electron affinity have larger electronegativity.

Group Trend

• The further down a group the farther the electron is away and the more electrons an atom has.

• So as you go from fluorine to chlorine to bromine and so on down the periodic table, the electrons are further away from the nucleus and better shielded from the nuclear charge and thus not as attracted to the nucleus. For that reason the electronegativity decreases as you go down the periodic table.

Period Trend

• Electronegativity increases from left to right across a period

• When the nuclear charge increases, so will the attraction that the atom has for electrons in its outermost energy level and that means the electronegativity will increase

Ionization energy, electronegativity

Electron affinity INCREASE

Atomic size increases,

shielding constant

Ionic size increases

3rd IE• Using the same logic s2p1 atoms have an low

3rd IE.

• Atoms in the aluminum family form +3 ions.

• 2nd IE and 3rd IE are always higher than 1st IE!!!

Electron Affinity

• The energy change associated with adding an electron to a gaseous atom.

• Easiest to add to group 7A.

• Gets them to full energy level.

• Increase from left to right atoms become smaller, with greater nuclear charge.

• Decrease as we go down a group.

Ionic Size

• Cations form by losing electrons.

• Cations are smaller than the atom they come from.

• Metals form cations.

• Cations of representative elements have noble gas configuration.

Ionic size

• Anions form by gaining electrons.

• Anions are bigger than the atom they come from.

• Nonmetals form anions.

• Anions of representative elements have noble gas configuration.

Configuration of Ions

• Ions always have noble gas configuration.

• Na is 1s22s22p63s1

• Forms a +1 ion : 1s22s22p6

• Same configuration as neon.

• Metals form ions with the configuration of the noble gas before them - they lose electrons.

Configuration of Ions

• Non-metals form ions by gaining electrons to achieve noble gas configuration.

• They end up with the configuration of the noble gas after them.

Periodic Trends

• Across the period nuclear charge increases so they get smaller.

• Energy level changes between anions and cations.

Li+1

Be+2

B+3

C+4

N-3

O-2 F-1

Size of Isoelectronic ions

• Iso - same

• Iso electronic ions have the same # of electrons

• Al+3 Mg+2 Na+1 Ne F-1 O-2 and N-3

• all have 10 electrons

• all have the configuration 1s12s22p6

Size of Isoelectronic ions

• Positvie ions have more protons so they are smaller.

Al+3

Mg+2

Na+1 Ne F-1 O-2 N-3

Lewis Dot Diagrams

• Atoms of different elements have different numbers of electrons

• Each shell is “full” before electrons move to the next shell.

• Period number: number of electron shells

• Shell will fill in the following order: 2, 8, 8, 18

• Group number: number of valence electrons (look at the second digit of the group number)

Trends:

• Elements of the same group have similar properties because they have the same number of valence e-

Noble gas configuration

• Group 18, the noble gases are the most stable of elements because their valence shell is full with electrons

• Stable octet

Lewis Dot Diagrams

• Visual representations of an element and only its valence e-

• Electrons get paired up along 4 sides of the element, max 2 e- per side.

• *Hydrogen will only have a max of 2 e-

Ionic Bonding (metal and nonmetal)

• Ionic bonding: attraction between oppositely charged ions formed when metallic ions (+) transfer electron(s) to nonmetallic ions (-)

• Ex: NaCl

Not always 1:1 ratio, sometimes need to use subscript to show the number of atoms

Ex: CaCl2 The 2 is a subscript, it shows that 2 atoms of chlorine bond with one atom of calcium.

The charges need to have a sum of zero

• A cation is a type of ion that has given away e-. It has a positive charge

Ex. Ca2+, K+

• An anion has accepted extra e-. It has a negative charge.

• Ex. F-, S2-

Covalent/molecular bonding

• Covalent bonds in nonmetals to create molecules

• Formed when a pair of electrons is shared by two atoms that are non-metals.

• A single bond is formed when 1 pair of e- is shared, double bond 2 pairs shared, triple bond 3 pairs shared

Single bonds

Draw the Lewis structures for the following molecules

(a) O2 (f) NH3

(b) H2 (g) H2O

(c)N2 (h) CO2

(d) F2

(e) Cl2

Write the equations (include e)

to show ionic bonds formed

between the following

(a) K and F

(b) Na and S

(c) Be and O

(d) Sr and N

Bonding

• Ionic Bonding - the atoms are so different that one or more electrons are transferred to form oppositely charged ions. The bonding results from the attractions between these ions.

• Covalent Bonding - two identical atoms share electrons equally. The bonding results from the mutual attraction of the two nuclei for the shared electrons.

Properties of Ionic and Covalent Bonding

• Properties of Ionic Compounds:

• They have relatively high melting points due to strong attraction in bonds

• They conduct electricity when molten or dissolved in water (ions can move freely)

• Those in solid state are not electrical conductors (ions cannot move)

• Ionic compounds dissolved in water form electrolytic solutions. Electrolyte – a substance that dissolves in water to produce a solution that conducts electricity.

• Properties of Covalent Compounds:

• They have relatively low melting points due to weaker attraction in bonds

• They tend not to conduct electricity when in solid or liquid state or dissolved in water (do not form ions; non-electrolytes)

Naming Ionic Compounds

• A compound can be identified either by its formula or it name. When naming ionic compounds, the following steps may be used:

1. Write the English name of the cation (metal)

2. Write the name of the anion by adding the suffix “ide” to the stem of the name of the non-metal from which it is derived.

Examples:

1. NaCl

2. KI

3. CaS

4. MgO

5. CaCl2

• This is the method used when transition metals are involved because they can have more than one type of cation. Follow these steps:

1. Write the English Name for the transition metal

2. Write the charge of the metal in roman numerals in round brackets (I, II, III, IV,..)

3. Write the non-metal by replacing the end of the name with the suffix “ide”

Examples:

1. FeCl3

2. FeO

3. Cu2S

4. PbO2

5. Fe2O3

6. HgO

Naming Molecular Compounds

• Follow these steps for naming:– The first word is the name of the element that

appears first in the formula. *A Greek prefix is used to show the number of atoms of that element in the formula.

– The second word consists of:• The appropriate Greek prefix designating the number

of atoms of the second element

• The root of the name of the second element

• The suffix “ide”

• Greek Prefixes:Number Prefix

1 Mono

2 Di

3 Tri

4 Tetra

5 Penta

6 Hexa

7 Hepta

8 Octa

9 Nona

10 Deca

Examples:

1. N2O5

2. CO2

3. NH3 (ammonia)

4. H2O2 (peroxide)

5. NO (nitric oxide)

6. N2O (nitrous oxide)

7. CH4 (methane)

Polyatomic Ions

• Polyatomic ions are ions that contain two or more atoms. They are electrically charged, which means that they are always looking at gaining or losing one or more electrons. So, polyatomic ions basically look like molecules (only contain non-metals), but with a charge.

• Although electrons can be lost/gained, most polyatomic ions are anions, meaning they have more electrons than the neutral atoms that are in the ion.

• So, hydroxide (OH-) has one extra electron beyond what the neutral oxygen and hydrogen atoms would have.

• The extra electron(s) do not belong to a particular atom. No matter how many extra electrons there are, they are shared equally among all atoms.

• Although most are anions, there are some, which are cations. The most common is NH4+ (ammonium).

Naming Compounds with Polyatomic Ions

• It is the same for naming ionic compounds but you must use the names of the polyatomic ions.

Examples:

1. LiOH

2. NH4Cl

3. BaSO3

4. Ba3(PO4)2

Naming binary compounds containing Hydrogen

• When hydrogen creates a compound with a metal it has different naming rules than when it creates a compound with a nonmetal.

H with a Metal.

- name of cation + hydride

Ex. CaH2

calcium hydride

H with a Nonmetal

• Hydrogen _____ide (no prefixes)

Ex. HCl

hydrogen chloride

Ex. H2S

hydrogen sulfide

Name!

1. KH

2. HI

3. H3N

4. MgH2