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Chemistry
chemical bond: force that holds two atoms together-creates stability in the atom
Bonds may form in two ways:1. Attraction between a positive nucleus and
negative electrons (covalent bonding)2. Attraction between a positive ion and a negative
ion (ionic bonding)
Remember: It is the valence electrons that are involved in this bonding.
ionic bond: electrostatic force that holds oppositely charged particles together
-called ionic compounds -forms between metals and nonmetals ◊metals lose electrons, forms a cation
~cation: positive ion from loss of electrons ◊nonmetals gain electrons, forms an anion
~anion: negative ion formed from gain of electrons
-most are binary, which means they contain 2 different
elements, such as MgO, Al2O3
Example:Sodium reacts with chlorine to form sodium
chloride.
Electron Configuration Notation:
Orbital Notation:
Lewis Dot Notation:
Try this # 1:Magnesium reacts with oxygen to form
magnesium oxide.
Electron Configuration Notation:
Orbital Notation:
Lewis Dot Notation:
Try this # 2:Lithium reacts with nitrogen to form lithium
nitride.
Electron Configuration Notation:
Orbital Notation:
Lewis Dot Notation:
It is the chemical bonds between atoms that determines many of the physical properties of the compound.
-alternating positive and negative ions form an ionic crystal
-the ratio of positive to negative ions is determined
by the number of electrons transferred -strong attraction results in a crystal lattice, a 3-D arrangement of atoms.
Other characteristics include:-high melting and boiling points-very hard and rigid-brittle-electrolyte when dissolved in water
During chemical reactions, energy is either absorbed (endergonic) or released (exergonic)
-the formation of ionic bonds is always exothermic (exergonic)
lattice energy: energy required to separate one mole of ions of an ionic compound
-the more negative the lattice energy, the stronger the bond
Lattice Energyies of Some Ionic Compounds
Compound Lattice Energy Compound Lattice EnergyName (kJ/mol) Name (kJ/mol)
KI -632 KF -808
KBr -671 AgCl -910
RbF -774 NaF -910
NaI -682 LiF -1030
NaBr -732 SrCl2 -2142
NaCl -769 MgO -3795
Depends on:1. smaller ions -more negative value because the
attraction is stronger between the nucleus and valence electrons
2. larger the positive/negative charge, the more negative the lattice energy because the attraction is stronger when more electrons are lost/gained
Lattice Energyies of Some Ionic CompoundsCompound Lattice Energy Compound Lattice Energy
Name (kJ/mol) Name (kJ/mol)KI -632 KF -808
KBr -671 AgCl -910RbF -774 NaF -910NaI -682 LiF -1030
NaBr -732 SrCl2 -2142NaCl -769 MgO -3795
A universal set of rules must be used so chemists around the world can communicate.
formula unit: simplest ratio of ions represented in an ionic compound
-remember that ionic compounds form a crystal lattice, consisting of many cations and anions.
-the overall charge for the compound is 0
Most ionic compounds are binary, consisting of two monatomic ions.
-monatomic ion: one atom ion, either positively or negatively charged
Remember that we determine the charge of each ion by its oxidation number.
Formula Rules for Ionic Compounds1. write the cation first, followed by the anion 2. state the charges of both ions3. cross the number for the charge of one ion to
become the subscript for the other ion. -subscripts are used to state the number of
each atom in the compound
Example: Determine the formula for the ionic compound formed when potassium reacts with oxygen.
1. Cation = potassium = K Anion = oxygen = O2. K+1 O-2
3. K+1 O-2
K2O1
K2O
You try: Determine the formula for the ionic compound formed when aluminum reacts with chlorine.
We write formulas for ionic compounds containing polyatomic ions the same way as in binary compounds.
-the cation comes first, followed by the anion -state the charges -cross over the number for the chargesHowever: -if you have more than one polyatomic ion, place parenthesis around the polyatomic ion, with the subscript outside the parenthesis.
Example: Determine the formula for the ionic compound formed when beryllium reacts with cyanide.
1. Cation = beryllium = Be Anion = cyanide = CN-
2. Be+2 CN-1
3. Be+2 CN-1
Be1(CN)2
Be(CN)2
You try: Determine the formula for the ionic compound formed when ammonium reacts with iodine.
The names of ionic compounds include the ions of which they are composed.
1. The element whose symbol appears first in the formula also appears first in the name.
-this is always the positively charged ion, or metal
2. The name of the second ion follows, with its ending
changed to –ide for single atom ions.
Ex: What is the name of MgCl2?
magnesium chloride
You follow the same rules when naming polyatomic ions as when you have binary ionic compounds, however:
-you do not change the ending of the polyatomic ions, even when they are the second atom.
Example: Al2(SO4)3
aluminum (III) sulfate
Rule: You must state the charge of all metals not included in groups 1 and 2 because many
have multiple charges.
*According to the previous rules, FeO and Fe2O3 would both be named iron oxide,even though they are not the same compound*
Since many transition metals can have more than one charge, the name must show this. This is done using roman numerals. -FeO is named iron (II) oxide because Fe has a +2
charge -Fe2O3 is named iron (III) oxide because Fe has
a +3 charge
*The roman numeral states the charge of the metal*
Q: How do I know the iron in FeO has a +2 charge? A: The oxide ion has a –2 charge, so the Fe must have a +2 charge so the compound is overall neutral.
Q: How do I know the iron in Fe2O3 has a +3 charge?
A: There are three oxide ions with a –2 charge:(3 ions)(-2 charge/ion) = a total of –6 charge
Since the overall charge must be neutral, the iron must have a total charge of +6. Therefore:
(2 ions)(x charge/ion) = +6 x = +3
Metallic bonds are similar to ionic bonds because they often form lattices in the solid state.
-eight to twelve metal atoms surround another, central metal atom
Instead of sharing electrons or losing electrons, the outer orbitals overlap.
-electron sea model: all metal atoms in a metallic
solid contribute their valence electrons to form a ‘sea’
of electrons around the metal atoms. -valence electrons are free to move from atom to atom (delocalized electrons), forming metallic cations
metallic bond: attraction of a metallic cation for the delocalized electrons that surround it
This bonding contributes to the unique properties of metals:
1. generally have high melting and boiling points, with
especially high boiling points -due to the amount of energy needed to separate the electrons from the group of cations 2. malleable (hammered into sheets) and 3. ductile (drawn into wire) -mobile electrons can easily be pulled and
pushed past each other
4. durable -though electrons move freely, they are strongly attracted to the metal cations and are not easily removed from the metal
5. good conductors -free movement of the delocalized electrons, allowing heat and electricity to move from
one place to another very quickly 6. luster -interaction between light and delocalized electrons
As the number of delocalized electrons increases, as in transition metals (d electrons), the hardness and strength also increases.
-alkali and alkaline earth metals are soft (s valence
electrons only)
It is easy to combine 2 or more different metals to make a metallic crystal
-alloy: mixture of elements with metallic properties
-the properties of alloys differ from those of the individual elements that make it up
TEST
Remember that atoms bond to increase stability, which occurs when an atom gets a full outer shell of electrons.
-in ionic bonding, one atom loses electrons (metal)
and another gains electrons (nonmetal) to form oppositely charged ions with a full outer shell
However, sometimes there is not a transfer of electrons, but a sharing of electrons.
-covalent bond: attractive force between atoms due
to the sharing of valence electrons
Covalent bonds can form between: -2 or more nonmetal atoms -metalloids (especially the ones to the right of the metalloid line) and nonmetals
molecule: when two or more atoms bond covalently
Covalent bonds can have either single bonds or multiple bonds.
-single bonds: 2 shared electrons (1 pair) -multiple bonds: 4 or 6 electrons shared (2 pair= double or 3 pair = triple)
When we show bonding, shared electron pairs can be shown by either a pair of dots or a single line.
-Lewis Structures are used to show how bonding electrons are arranged in molecules
-example: NH3
-sigma bond (): single covalent bond formed when
an electron pair is shared by the direct overlap of orbitals ♦can occur between s & s, s & p , or p & p orbitals
A multiple bond forms when two atoms share more than 2 electrons.
-double bond: 4 electrons shared ( 2 pairs) ♦ O2
-triple bond: 6 electrons shared (3 pairs) ♦ N2
Some molecules have both single and multiple bonds.
♦HCNpi bond (): forms when parallel orbitals overlap to
share electrons -only occurs with multiple bonds because the first overlap is always a sigma bond
All bonds can be broken, though some more easily than others.
-due to the strength of the bond
What affects bond strength?bond length: distance that separates the bonded
nuclei -determined by the size of the atoms and how many electron pairs are shared ♦larger the atom, the longer the bond length, the weaker the bond ♦more shared electrons gives a shorter, stronger
bond
When a bond forms or breaks, an energy change occurs.
-bond formation: energy released (exergonic) -bond breaking: energy absorbed (endergonic)
bond dissociation energy: amount of energy required to break a specific covalent bond
-always a positive number -indicates the strength of a covalent bond
larger the bond dissociation energy, stronger the bond
(see p 246 for examples)
1. low melting and boiling points.2. many vaporize readily at room temperature3. relatively soft solids (but not all, some are
gases/liq.)4. can form weak crystal lattices5. do not conduct electricity when dissolved in
water
These properties are due as a result of differences in attractive forces
-attraction between atoms within a molecules is strong
-attraction between different molecules is weak ~called intermolecular forces or van der Walls
forces
Types of Intermolecular Forces (van der Walls forces)1. dispersion force (induced dipole)2. dipole-dipole force3. hydrogen bonding
dispersion force (induced dipole) -occurs between nonpolar molecules -very weakdipole-dipole force
-occurs between polar molecules -the more polar the molecule, the stronger the
forcehydrogen bonding
-strong intermolecular force between the hydrogen end of one dipole and a fluorine, oxygen or nitrogen atom on another molecule’s dipole
Molecules are represented by both names and formulas.
Rules for Naming Binary Molecular Compounds1. The first element in the formula is named first, using the entire element name.2. The second element in the formula is named using the root of the element and adding the suffix –ide.3. Prefixes are used to indicate the number of atoms of each type that are present in the compound. -exception: 1st element never uses the prefix mono- -drop the final letter of the prefix if element name begins with a vowel.
Prefixes you need to know:# atoms prefix
1 mono- 2 di-
3 tri- 4 tetra-
5 penta- 6 hexa-
7 hepta- 8 octa- 9 nona- 10 deca-
Name the compound P2O5, which is used as a drying and dehydrating agent.
1st atom: P = phosphorus 2nd atom: O = oxygen = oxide
There are 2 phosphorus = diphosphorusThere are 5 oxygens = pentoxide (drop the –a of
penta-)
Put it together: diphosphorus pentoxide
(We will talk more about acids in Ch 19)There are two types of acids:1. binary acid: contains hydrogen and one other element -when naming use the prefix hydro- plus the root
of the second element with the suffix –ic, followed
by the word acid. -ex: HCl
H = hydro- Cl = chloride = chloric hydrochloric acid
Some acids are not binary, but are named according to the binary acid rules when oxygen is not present, as in HCN.
H = hydro CN = cyanide = cyanic
hydrocyanic acid2. oxyacid: an acid that contains an oxyanion
(oxygen containing polyatomic ion) -the name depends on the oxyanion present
-the name consists of the root of the anion, a suffix,
and the word acid ♦if the anion suffix is –ate, it is replaced with -ic ♦if the anion suffix is –ite, it is replaced with -
ous
-examples: ~HNO3
NO3 = nitrate
= nitric nitric acid ~HNO2
NO2 = nitrite
= nitrous nitrous acid
Use the prefixes in the molecule’s name to determine the subscript for each atom in the compound.
- phosphorus tribromide P Br
1 (no prefix) 3 (tri) PBr3
- the formula for an acid can be derived from the name as well
♦charge of the oxyanion or anion gives the number
of hydrogens hydrofluoric acid = HF
(1 H because fluorine has a -1 charge)
structural formula: uses letter symbols and bonds to show relative positions of atoms
-one of the most useful -can be predicted for many molecules by
drawing Lewis structures -H is always an end (terminal) atom, never a
central atom -less electronegative atom is the central atom
(nm or metalloid closest to the left of the PT-usually)
CH2O
1. Predict the location of the atomsC is least electronegative & farthest to left
on PT, therefore it is the central atom2. Find the total number of electrons available for bonding.
1 C-4, 2 H-2, 1 O-6 for a total of 12 valence e-
3. Determine the number of bonding pairs12 valence e- / 2 = 6 electron pairs
4. Place one bonding pair (single bond) between the
central atom and each terminal atom.
H C O
H5. Subtract the number of pairs you used in step
4 from the number of bonding pairs determined in
step 3.6 – 3 used = 3 e- pairs left
5. Subtract the number of pairs you used in step 4 from
the number of bonding pairs determined in step 3.
-take the remaining electron pairs and place electron
pairs around the terminal atoms to satisfy the octet
ruleH C O
H
6. If the central atom is not surrounded by 4 electron
pairs, it does not have an octet -convert one or two of the lone pairs on a
terminal atom to a double or triple bond between that
terminal atom and the central atom
H C O H
Practice:1. CH3Cl 2. NBr5
Writing structural formulas for polyatomic ions is the same with one exception:
-the total number of electrons may differ due to the
negative and positive charge. ♦negative charge, more electrons are present
SO4-2 add two electrons
♦positive charge, less electrons are presentNH4
+1 subtract one electron
Let’s look at CO3-2.
-when one or more valid Lewis structure can be
written for a molecule, resonance occurs -let’s look at another resonance molecule/ion:
NO3-1
-each molecule/ion that undergoes resonance behaves as if it only has one Lewis structure
Some molecules do not obey the octet rule.
Three reasons exist:1. when a small group of molecules have an odd
number of valence electrons: -NO2 for a total of 17 valance electrons-one
unpaired electron on N
2. Some form with fewer than eight, though this is relatively rare:
-B in BH3 is stable with six because it only has 3 valence electrons.
3. When the central atom has more than 8 electrons, which is referred to as an expanded octet.
-can occur in elements that are found in period three
or higher elements (because of the d orbitals). -P in PCl5
(1 s orbital, 3 p orbitals, and 1 d orbital)
TEST/QUIZ
Many of the physical and chemical properties of molecules is determined by the shape of the molecule.
-the shape of molecules determines if two or more molecules can get close enough for a reaction to occur.
VSEPR (Valence Shell Electron Pair Repulsion) model: atoms in a molecule are arranged so that the pairs of electrons (bonded and lone) minimize repulsion.
The repulsion between electron pairs result in fixed angles between atoms
-bond angle: angle formed by any two terminal atoms and the central atom
♦lone pairs take up slightly more space than bonded
pairs ♦multiple bonds have no affect on the geometry because they exist in the same region as single bonds -example: H2O
See page 260 for the Molecular Geometries (Shapes)
Remember that atoms have different attractions for electrons (electronegativity).
-electronegativity increases left to right and decreases
down a period
The character and type of bond can be predicted using the difference in electronegativities between bonded atoms.
-pure covalent bond: electronegativity difference = 0
(usually occurs between identical atoms, H2)
Most atoms do not have equal sharing of electrons, producing a purely covalent bond.
-polar covalent bond: unequal sharing of electrons ♦the larger the electronegativity difference, the
more ionic the bond character -ionic bonds form when the electronegativity difference is > 1.7 and nonpolar covalent bonds
form when the difference is < 0.5 -the cutoff between polar covalent, nonpolar, and ionic is sometimes inconsistent with experimental data
Remember: bonding is not clearly ionic or covalent, with ionic character increasing as the difference in electronegativity increases.
Decide if the following pairs of atoms are polar covalent, nonpolar covalent or ionic.
1. N-H3.04-2.20 = 0.84
polar covalent2. C-Cl 2.55-3.16 = 0.61 polar covalent3. S-Se
2.58-2.55 = 0.03 nonpolar covalent
When a polar bond forms the shared electrons are pulled more strongly toward one atom.
-this creates partial charges at opposite ends of the molecule, which is called a dipole
♦ - indicates a partial negative + indicates a partial positive
Polar molecule or not?A molecule can have individual polar bonds, but
make a nonpolar molecule. How?We look at the shape of the molecule.
Let’s look at H2O and CCl4.
O—H C—Cl - + + -
1.24 0.61both O-H and C-Cl have polar covalent
bonds
One molecule is polar and the other is nonpolar? How do we know?
We look at the shape of the molecule and the terminal atoms.
-symmetric molecules like CCl4 are nonpolar because the polar bonds cancel each other out.
CCl4
-asymmetric molecules like H2O are polar because the polar bonds do not cancel each other out.
H2O
If water is polar, why will oil not dissolve in it?Oil must be nonpolar because
A substance is only soluble (dissolvable) when combined with a like molecule.
“Like Dissolves Like”
hydrophobic- “fear of water”hydrophilic- “likes water”
TEST