Chapter 9 - Chemical Bonds
Last chapter - atoms:smallest unit of an element
-properties determined by electron configuration
-combine together to form compounds
compound - substances made up of
combinations of elements
compounds formed or broken down in
chemical reactions - changes in matter
when different chemical substances are
created by forming or breaking
“chemical bonds”-bind atoms together
H2 and O2 (diatomic)
form H20
new properties due to
changes in e- configuration
Remember noble gases (stable)-never combine
monatomic molecules - gases
molecules - special bonding arrangements
acting as independent units - neutral groups of atoms
chemical bonding between electron orbitals
cause substances to change their properties
Chemical bonding involves changes in the
internal potential energy - chemical energy
Compounds absorb energy to increase chemical energy
-release energy to lower chemical energy
Example:PHOTOSYNTHESIS
sunlight+carbon dioxide+water
plant material+oxygen
{CHEMICAL EQUATIOMN- blueprint for reaction
-read: reactants yield products}
plants increase chemical energy - store radiant energy
can also release energy - burning of wood
plant + oxygen
carbon dioxide+water+energy
releases all of the sunlight energy stored
-energy stored by forming bond
-energy released by breaking bond
The energy exchanges are achieved
through changes in the electronic structure
of the atoms
Valence electrons - outer shell electrons
of an atom
- highest n value
- usually s & p orbitals (1 to 8 electrons)
- “fingers of atom”
- determine chemical properties of an element
- valence = group number
Representative (A group) elements attempt to
acquire an outer orbital with
eight electrons
octet rule
(except H - s2 orbital stable)
rule of thumb: metals lose electrons (+)
non-metals gain electrons (-)
want noble gas configuration
closed “fingers”
Example: how sodium form stable configuration?
Na Group IA - 1 valence electron (metal-lose e-)
.
electron dot: Na + energy
Na+ + electron
What are the charges for the following ions:
Ca, Br, Al, O, P
Types of Chemical bonds
chemical bond - attractive force that holds atoms
together in a compound
-atoms gain or lose electrons to
achieve lower energy
-most stable arrangement “octet”
Three types of bonding:
ionic bond- atoms transfer electrons
-Coulombic attraction between ions
-generally between metals & non-metals
covalent bond- atoms share electrons
-share valence electrons to form octet
-generally bonding between non-metals
metallic bond- all of the metal atoms
share electron sea
-gives metallic properties
- we will not consider this bonding
IONIC BOND - transfer of electrons
chemical bond of electrostatic attraction - Coulomb’s Law
valence electrons are lost or gained to achieve
stable noble gas configuration - octet
form solids-water soluble as ions
BOND BETWEEN METALS AND NON-METALS
metals lose electrons
positive ions
non-metals gain electrons
negative ions
opposite charged ions attract
energy is released when bond forms
lower chemical energy
heat of formation - energy released when an
ionic bond is formed
-same energy required to break bond
energy conserved in decomposition
. .......
Illustration:
Na + Cl
Na+ Cl- + energy
Na+ formation requires energy
Cl- formation releases energy
Cubic
net energy released
crystal
forms crystalline solid - each ion surrounded
by opposite ion in one giant unit made of charged ions
Ionic Bonds continued
remember electrons transferred to form octet:
isolated atoms with extra or less charge
-electroneutrality electrons gained by one atom
equals electrons lost by other
Describing IONIC COMPOUNDS: held by ionic bonds
chemical formula - describes what elements
are present and in what proportions
-uses element symbols
-and subscripts for number
of atoms in compound
-no subscript if only one atom
NaCl understood 1Na & 1Cl
must know charge on atoms:
Group IA and IIA- positive charge = family
Group VI & VII-negavive charge=8-family
Transition metals - form positive ions
variable charge - several possible
charge values
i.e., copper
Cu+ and Cu+2
Ionic chemical formula determined by
electroneutrality
Charge on positive ion(s) must balance charge on negative ions
Example: what is the chemical formula for
the ionic compound containing calcium
and fluorine?
Ca:Group IIA - positive ion (metal)
loses 2 e- to form octet (Ar config)
Ca+2
F: Group VIIA - negative ion (non-metal)
gains 1 e- to form octet (Ar again)
FNEED 2F- to cancel 1Ca+2
alternatively, 2 e- from Ca tranferred to 2F
chemical formula : CaF2
Transition metals- Table 9.2 different charges possible
-form colored solids
-multiple chemical formulae with
same elements
Cobalt with Chlorine (Cl-)
Co+2
Co+3
CoCl2
CoCl3
form different compounds with different properties
Polyatomic Ions
Groups of two or more elements (poly) that form charges
bond together covalently - share electrons
separate unit with net charge
that act like single ions
HAVE SPECIFIC NAMES AND CHARGES
Table 9.3
sulfate (SO4-2)
nitrate (NO3-)
typically bond with metals like Ca:
CaSO4
Ca(NO3)2
( ) needed if more than
one needed fo rbond
Covalent Bond - chemical bond formed by the
sharing of an electron pair
-tightly bound in a molecule - smallest unit of
covalent compound
-electrons move throughout entire molecular orbital
-both elements count the electron pair
to achieve octet
-bonding between NON-METALS
Halogens GroupVII (non-metals)
reactive elements that often form diatomic molecules
Example of covalent bond: fluorine
...F...F
xx
x
x
Halogen gas
both electrons counts as octet
x x
nonmetals
F-F
for each atom - shared
represents
covalent bond - dash is bonding pair - shared
-keeps molecule
together
orbital pairs: - two electrons (s=+/-1/2) in orbital
lone pairs that are not shared in bond
x
WATER H2O
...O...
. .
H
H
O -- octet
O
H
H
H -- He config
SINGLE COVALENT BOND - single electron pair
shared by each atom
DOUBLE COVALENT BOND - two electron pairs
shared by each atom
each element must have 2 unpaired e-: C, N, O, S
Example: ehtylene (gas from ripening fruit ) C2H4
Hx
H
H
H
double dashx
C
Cx
C C
double bond
x
H
H
H
H
TRIPLE BOND - 3 pairs of electrons shared (C,N) H C C H
... .
. .
ACETYLENE
Ionic versus Covalent
General rule:
ionic - opposite sides of periodic table
(metal & non-metal
covalent- bonding between non-metals
Depends on how tightly elements can hold onto
their outer electrons
compounds can have ionic and covalent properties
Electronegativity - comparative ability of atoms to
attract bonding electrons FIG. 9.10
- higher number, attracts electrons well
-will pull off e;ectrons better to form ionic
-about same number
share electrons
covalent
ABSOLUTE ELECTRONEGATIVITY DIFFERENCE
subtract electronegativity of two elements-predict bond type
difference>1.7
difference <0.5
steals electrons - ionic
shares electrons - covalent
Polar molecule between 0.5 and 1.7
- unequal sharing
one side of molecule more (+)
other side more (-)
Electronegativity
Table 9.10
F holds onto electrons best - largest
electronegativity
Fr has smallest electronegativity - easily gives
up electrons
mixed polar molecules not charged
-just unequal sharing
Examples:
HF
-
+
difference=
ionic
separate charges
N2
difference=
covalent
share equally - diatomic
HCl
difference=
polar covalent
+
-
Electronegativity difference describes percentage ionic
how close it is to removing electrons
How to name compounds
Historically: name based on usage
“soda” sodium carbonate (from sodium)
baking soda-sodium bicarbonate
caustic soda (burns)-sodium hydroxide
pot ash – potassium carbonate
washing soda, etc.
STANDARDIZE: so everyone knows what compound
systematic rules -can determine composition
RULE: if two elements, ends in –ide (unless
polyatomic ion)
Naming different for ionic and covalent
IONIC COMPOUND NAMES
Representative Elements (A groups)
metal first, then non-metal with –ide ending
NaCl sodium chloride
MgF2 magnesium fluoride
definite charge – only one compound possible
Transition Metals (B groups)
variable charges – different compounds possible
need to indicate chargebalance charges to get formula
old way- latin root: -ic higher charge state
-ous lower charge state
oxidation number -charge on ion
Ionic compound names – transition metals
oxidation numbers: Table 9.6
Modern way to indicate oxidation state – Roman
numerals
Roman numeral in parenthesis- (II) means +2
Fe+3 iron (III)
Fe+2 iron (II)
(Fe2)+6(S3)-6 iron(III) sulfide
Name these:
What is the formula?
CuCl
lead(II)sulfide
SnO
chromium(III)bromide
Fe2S3
lead(II)oxide
Ionic compounds with polyatomic ions:
two or more elements-act like single ion
-have specific names Table 9.3
-ate: one more O
hydroxide (OH)than –ite
-2
sulfate (SO4)
-per: one more O
than –ate
nitrate (NO3)-hypo: one less O
than -ite
Name the following:
Na2SO4
Ca(NO3)2
Fe(NO3)2
write formula:
rubidium chlorate
lead(II) sulfate
barium permanganate
Summary: Method for writing ionic formula
1. Write positive element first (metal) then least metallic
2. Balance number of ions (polyatomics) by making neutral
cross-over technique: decide numbers of each ion
in compound by using charge number of
the other element (or polyatomic)
Ca+2 & Cl-1
Sn+2 & F-1
Mg+2 & O-2
iron(III) oxide :
lead(II) phosphate :
zinc carbonate :
magnesium hydroxide :
CaCl2
SnF2
MgO
Covalent Compound Names
Molecules – made of non-metals
Problems:
-no metal to put first!
-possibility of multiple bonds
two elements form various
compounds
Name must indicate compound better!!!
RULES: for two element covalent compounds
1. Least metallic (to the left) element first with prefix
if number more than one atom
2. Stem name of 2nd element (with –ide ending)
adding prefix only if elements form more
than one compound
-only one C
Examples: CO
carbon monoxide
-multiple compounds
CO2
carbon dioxide possible
Formula tells you
how to write name
Name:
N 2O 2
N 2O 3
CS2
dinitrogen dioxide
dinitrogen trioxide
N 2O
laughing gas
CCl4
Covalent name tells you how to write formula
EXAMPLES: write chemical formula for the following:
manganese dioxide
dinitrogen pentoxide
dihydrogen dioxide
this very specific naming necessary
because you need to know how they bond
in order to get the chemical formula
O=C=O
CO2 can form double bonds
-but C will always form same
number of covalent bonds
-combinations of single,
double and triple
same number of bonding pairs
VALENCE – number of covalent bonds an atom can form
Carbon
....
C
H valence of 1
O valence of 2
N valence of 3
4 bonding electrons
valence of 4 – can form 4 covalent bonds
4 single, 2 double,
1 double - 2 single, etc.