Valence electrons

Electrons in the highest occupied energy level of an atom.

These are the electrons that determine the element’s properties.

Electron dot structures Diagrams that show the

valence electrons as dots. The core electrons and

the nucleus are included in the symbol of the element

Octet Rule

In forming compounds, atoms tend to achieve noble gas configuration.

8 electrons in the highest energy level.

IONS Atoms or groups of atoms that

have a positive or negative charge.

Cations - positive ion resulting from loss of electrons.

Anions - negative ions resulting from gain of electrons.

Formula Unit

Represents an ionic compound.

Lowest whole number ratio of ions in the compound.

Formula units Ionic compounds form as

repeating links in a crystal matrix.

Each cation is bound to each neighboring anion

The ions are “locked in place”

NaCl is the smallest ratio that indicates this

matrix

Ionic Bond

Bonds resulting from the electrostatic attraction between oppositely charged ions. In an ionic compound the net ionic charge is 0.

Ionic Compounds

Metal + Non-metal Polyatomic cation+Non-metal Metal+ Polyatomic anion Solid at room temperature High melting point >300°C

Ionic compounds NaCl Na2SO4

CaCO3

1. Crystalline solids that have high melting points.

2. They are often soluble in water

3. They conduct electricity when in solution, or when molten.

Predicting the formula1. Identify the charge of

the cation2. Identify the charge of

the anion3. Make a “T” table4. Add ions until the

positive charge equals the negative charge

Predicting the formula

Polyatomic ions are groups of atoms that stay together, they are treated like “super atoms” The entire group carries the charge.

Table 9.3 p.257

NH

4

+

NH

4

+

Lets practice !!!!Potassium + Phosphorus

Lithium + Selenium

Aluminum + Chlorine

Gallium + Sulfur

Magnesium + Iodine

Sodium + Carbonate

Sodium + Hydrogen Carbonate

Strontium + Phosphate

Ammonium + Chromate

Barium + Acetate

Lets practice !!!!

Potassium + Phosphorus K3P

Lithium + Selenium Li2Se

Aluminum + Chlorine AlCl3Gallium + Sulfur Ga2S3

Magnesium + Iodine MgI2Sodium + Carbonate Na2CO3

Sodium + Hydrogen Carbonate NaHCO3

Strontium + Phosphate Sr3(PO4)2

Ammonium + Chromate (NH4)2CrO4

Barium + Acetate Ba(C2H3O2)2

Lets Review!!!calcium chloride

cesium oxide

aluminum perchlorate

barium sulfide

sodium dichromate

aluminum phosphate

calcium carbonate

sodium carbonate

Lets Review!!!calcium chloride CaCl2cesium oxide Cs2O

aluminum perchlorate

Al (ClO4)3

barium sulfide BaS

sodium dichromate Na2Cr2O7

aluminum phosphate

Al PO4

calcium carbonate CaCO3

sodium carbonate Na2CO3

Compounds with transition metals Transition metals can

have more than one charge.

You may have more than one possible compound: FeO, or Fe2O3

Make tables & work backwards to determine cation charge

Indicate charge with a roman numeral

REMEMBER THE TABLE MUST BE BALANCED !!!!!

Fe+? O-2

Fe+? O-2

Fe+? O-2

O-2

iron(II) oxide

iron(III) oxide

Some Ions we need to just Know

Silver is always +1 Zinc is always +2 Cadmium is always +2 Do not use a roman numeral with these Iron may be +2 or +3 Tin may be +2 or +4 Lead may be +2 or +4 More in table 9.2 p.255

Lets Review!!!potassium oxide

strontium nitride

strontium nitrate

strontium nitrite

aluminum hydroxide

magnesium sulfate

iron(III) oxide

silver oxide

Lets Review!!!

Potassium Oxide K2O

Strontium Nitride Sr3N2

Strontium Nitrate Sr(NO3)2

Strontium Nitrite Sr(NO2)2

Aluminum Hydroxide Al(OH)3

Magnesium Sulfate MgSO4

Iron(III) Oxide Fe2O3

Silver Oxide Ag2O

Lets practice!!

Na2S

Hg2S

Na2Cr2O7

Hg2Cr2O7

CuO

Lets Practice Answers

Na2S sodium sulfide

Hg2S mercury(I) sulfide

Na2Cr2O7 sodium dichromate

Hg2Cr2O7 mercury(I) dichromate

CuO copper(II) oxide

Lets Practicecalcium carbonate

ammonium sulfate

copper(I) phosphate

chromium(IV) acetate

cadmium perchlorate

Lets Practice

calcium carbonate CaCO3

ammonium sulfate (NH4)2SO4

copper(I) phosphate Cu3PO4

chromium(IV) acetate

Cr(C2H3O2)4

cadmium perchlorate

Cd(ClO4)2

Links to practice tests and games

Interactive link Interactive link 2 Interactive link 3 Interactive link 4

Metallic Bonds The force of attraction that

holds metals together. The attraction of the free floating electrons for the positively charged metal ions

Metallic Properties

Malleable Ductile Conduct heat and

electricity.

Single covalent bond

A bond in which two atoms share a pair of electrons between them in order to achieve noble gas configuration.

Structural formulas

Chemical formulas that show the arrangement of atoms in molecules and polyatomic ions. Each dash represents a pair of shared electrons.

Unshared pairs Pairs of valence electrons

that are not involved in bonding, not shared between atoms.

Also called lone pairs or non-bonding pairs

Double covalent bond

Two atoms share two pairs of electrons between them to attain noble gas configuration.

O2 and CO2

Triple covalent bond

Two atoms share three pairs of electrons between them to attain noble gas configuration.

N2

Coordinate covalent bond

A covalent bond in which one atom contributes both bonding electrons.

CO and NH4+ and N2O

Exceptions to the octet rule

NO2

BF3

PCl5

SF6

Law of Definite Proportions

In samples of any chemical compound, the masses of the elements are always in the same proportions.

Law of Multiple Proportions When two elements form

more than one compound, the different masses of one element that combine with the same mass of the other element are in the ratio of small whole numbers.

Isoelectronic ions Ions containing the same

numbers of electrons. Generally for

isoelectronic ions size decreases as nuclear charge increases.

Bond energy The energy required to

break a bond. Table 8.4 p.365 Bond length- the distance

where energy is minimum.Table 8.5 p.365

Coulomb’s Law E=2.31x10-19Jxnm Q1Q2

r E= energy in joules r= distance between ion

centers in nm Q1&Q2= ion charges

Molecular Compounds All non-metals covalently

bonded. Solid, liquid or gas Low melting point <300°C Smallest representative

particle is a molecule.

Non-polar covalent bond A covalent bond in

which the electrons are shared equally. The two atoms have nearly the same electronegativities

Polar covalent bond A covalent bond in which

the electrons are not shared equally. The more electronegative atom will pull more of the electrons toward itself.

Polar molecule One end of the molecule

has a slightly positive charge and one has a slightly negative charge.

This is called a dipole. Depends on the shape.

Lattice energy The change in energy that

takes place when separated gaseous ions are packed together to form an ionic solid.

Lattice Energy=k(Q1Q2)

r

Use the following to calculate H°f of BaCl2(s).

Lattice energy= -2056 kJ/mol 1st ionization Ba= 503kJ/mol 2nd ionization Ba= 965kJ/mol Electron affinity Cl=-348kJ/mol Bond energy Cl2=239kJ/mol

H sublimation Ba=178kJ/mol

Bond Energies & Enthalpy

H=D(bonds broken)- D(bonds formed) =sum of terms D=bond energy per mol of

bonds, always positive.

Localized Electron Bonding Model A molecule is composed of

atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms.

Lewis Structures VSEPR Theory

Writing Lewis Structures Sum the valence electrons

from all the atoms. Use a line to show a pair of

electrons between each pair of bound atoms. (Bonding Pairs)

Writing Lewis Structures Arrange the remaining

electrons to satisfy the duet rule for hydrogen and the octet rule for the 2nd row elements. (Lone Pairs)

Double or triple bonds may be needed.

Comments on Octet Rule C,N,O,F obey octet rule. B and Be often have less than

8 electrons. Very reactive. 2nd row never exceed rule. 3rd row and up often obey

octet rule but may exceed it., due to d orbitals.

When writing lewis structures satisfy the octet rule for the atoms first. Place any remaining electrons on the elements that have available d orbitals.

Resonance structures Occur when it is possible

to have two or more valid electron dot structures for the same molecule or ion.

SO3, SO2

Formal Charge (FC) A method to decide which

of many possible non-equivalent Lewis structures is most likely to occur.

Atoms in molecules try to achieve FC as close to 0 as possible.

FC=(# valence e- on free atom) -(# valence e- assigned to the atom in the molecule).

(Valence e-)assigned = (# lone pair electrons) + 1/2(#shared electrons)

VSEPR theory Valence-shell-electron-pair

repulsion theory. Because electron pairs repel

molecular shape adjusts so the valence electron pairs are as far apart as possible.

Hybrid orbitals

In hybridization several atomic orbitals mix to form the same number of equivalent hybrid orbitals

Sigma bonds

Formed along the axis that joins the atomic nuclei when two atomic nuclei combine to form a molecular orbital.

Pi bond

Electron in pi bonds are found in sausage shaped regions above and below, or in front and behind the bond axis.

Paramagnetic molecules

Show an attraction to an external magnetic field.

Molecules contain one or more unpaired electrons.

Diamagnetic Molecules

Molecule is repelled by an external magnetic force.

Associated with paired electrons

Electronegativity The ability of an atom in a

compound to draw electrons to itself.

Pauling electonegativity values Table 14.2 p.405

Large electronegativity differences correspond to ionic bonds