Chapter 5
(section 1)
Chemical bonds distinguish compounds from mixtures
A compound always has the same chemical formula
Always made of same elements in the same proportion
Chemical structure shows the bonding within a compound

bond length gives distance between the nuclei of two bonded atoms

bond angles tell how these atoms are oriented in space
Some models give you an idea of bond lengths & angles ball and stick
Space-filling models show the space occupied by atoms ball and attached half balls

Compounds with network structures are strong solids
Some compounds are made of networks of bonded ions
Some compounds are made of molecules
The strength of attractions between molecules varies
Attraction between water molecules are called hydrogen bonds strong bonds within each water molecule weaker attractions between water molecules
(section 2)
three different kinds of bonds
Bonded atoms usually have a stable electron configuration full p & s orbitals more stable than partly filled p & s orbitals

Bonds can bend & stretch without breaking
formed between oppositely charged ions
Ionic bonds are formed by the transfer of electrons not sharing to fill outermost energy levels, transfer electrons instead
Ionic compounds are in the form of networks, not molecules ratio formula unit
When melted or dissolved in water, ionic compounds conduct electricity
 solid ionic compounds do not conduct electricity because ions are locked into place
 liquid ionic compounds conduct electricity because ions are free to move
Electrons move freely between atoms
 metallic bond- bond formed by the attraction between positively charged metal ions
& the electrons around them
 electrons free to move from atom to atom (conducts electricity)

 solid, liquid or gas
 most do NOT conduct electricity
Atoms joined by covalent bonds share electrons
 no transferring because both need to gain electron(s)
 sharing of the valence electrons

nonpolar covalent bonds-sharing equally
Atoms may share more than one pair of electrons
 double bonds- two pairs of electrons is shared (total 4)
 triple bonds- three pairs of electrons shared (total 6) examples:
O=O
NΞN
 triple bonds stronger than double bonds because it requires more energy to break triple bonds
Atoms do not always share electrons equally

polar covalent bonds- not sharing equally (unequal sharing)
 usually electrons are more attracted to atoms of elements that are located to the further right & close to top of periodic table

Covalently bonded atoms that have gained or lost electrons
There are many common polyatomic ions examples:
 baking soda: sodium hydrogen carbonate, NaHCO
3 found in toothpastes used in baking
 sodium bicarbonate, Na
2
CO
3 used to make soap & cleaners
Parentheses group the atoms of a polyatomic ion example:
 ammonium sulfate, (NH
4
)
2
SO
4 parentheses are to remind you that the atoms of the ammonium ion act as a single ion
Some polyatomic anion names relate to their oxygen content
-ite, less
-ate, more

(section 3)
Names of cations include the elements of which they are composed common cations: cesium, Cs + lithium, Li + potassium, K + rubidium, Rb + sodium, Na + barium, Ba 2+ beryllium, Be 2+ calcium, Ca 2+ magnesium, Mg 2+ strontium, Sr 2+ aluminum, Al 3+

Names of anions are altered names of elements fluorine, F become s fluoride, F - chlorine, Cl becomes chloride, Cl bromine, Br becomes bromide, Br iodine, I becomes iodide, I oxygen, O becomes oxide, O 2sulfur, S becomes sulfide, S 2nitrogen, N becomes nitride, N 3-
Some cation names must show their charge copper (I), Cu + copper (II), Cu 2+ iron (II), Fe 2+ iron (III), Fe 3+ nickel (II), Ni 2+ nickel (III), Ni 3+ chromium (II), Cr 2+ chromium (III), Cr 3+ cadmium (II), Cd 2+ titanium (II), Ti 2+ titanium (III), Ti 3+ titanium (IV), Ti 4+
Determining the charge of a transition metal cation

 ionic compounds have a total charge of zero
 total positive charges must equal total negative charges
named differently than ionic compounds
Numerical prefixes are used to name covalent compounds of two elements
1 mono
2 di
3 tri
4 tetra
5 penta
6 hexa
7 hepta
8 octa
9 nona
10 deca
A compound’s simplest formula is its empirical formula atoms in the simplest ratio example:
H
2
0 Hydrogen to oxygen is 2:1

Determining empirical formulas
142 g of unknown sample contains phosphorus, P & oxygen, O
P: 62 g P x 1 mol P / 30.97 g P = 2.0 mol P
O: 80 g O x 1 mol O / 16.00 g O = 5.0 mol 0 empirical formula is P
2
O
5 ratio 2:5 found to be 62 g P & 80 g O
Different compounds can have the same empirical formula
this is possible because empirical formula ONLY represents ratio example: compound empirical formula molar mass molecular formula formaldehyde CH
2
O acetic acid CH
2
O glucose CH
2
O
30.03 g/mol
60.06 g/mol
180.18 g/mol
CH
2
O
2 x CH
2
O = C
2
H
4
O
2
6 x CH
2
O = C
6
H
12
O
6
Molecular formulas are determined from empirical formulas
 not a ratio
 tells you how many atoms are in one molecule
 sometimes they are the same example: H
2
O