chapter four: forces between particles
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Chapter Four: Forces Between Particles. 2, 12, 14, 20, 22, 26-32, 36, 38, 48-58, 62, 66-74. Chemical Bonding Review. Compounds and Molecules are held together by chemical bonds Three types of bonds Ionic Metals and non-metals Covalent Non-metal and Non-metal Metallic - PowerPoint PPT PresentationTRANSCRIPT
Chapter Four: Forces Between Particles
2, 12, 14, 20, 22, 26-32, 36, 38, 48-58,
62, 66-74
Chemical Bonding Review• Compounds and Molecules are held
together by chemical bonds• Three types of bonds
– Ionic• Metals and non-metals
– Covalent• Non-metal and Non-metal
– Metallic• Between atoms of metals
Octet Rule• All atoms strive to have electronic
configurations like the Noble Gases• Eight electrons in the outermost shell,
highest principle quantum number (n)• Except H and He follow duet rule
– Want two electrons in outermost shell
• How do the atoms achieve an octet?
Taking or Giving and Sharing Electrons
• Ionic Bonds– Atoms take or give electrons from other atoms
• Covalent Bonds– Atoms share electrons between themselves
• Metallic Bonds– Sea of electrons
n=1
87654321
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87654321
87654321
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87654321
1
n=2
n=3
n=4
n=5
n=6
n=7
Valence Electron Review2
Lewis Dot Structures For Atoms/Ions
• Symbol represents the nucleus and all electrons except for those in the valence shell
• Give the Lewis Dot Structure for:Na F O2-
• Species with the same number of electrons are isoelectricO2- F- Ne Na+
Mg2+
How many electrons does each species have?
Lewis Dot Structures• GN Lewis developed the theory of
covalent bonding• Structures showing covalent bonds are
called Lewis structures• Each line represents a shared pair of
electrons (2 electrons)• Lone pairs of electrons are shown by a
pair of dots
Drawing Lewis Structures• Decide on atom connectivity and placement
– Hydrogen (never in the middle) is frequently bonded to oxygen
– Oxygen is rarely the central atom– Oxygen will not bond to oxygen (except O2 or O3)– Carbon will be the central atom– Least electronegative atom is in the middle
Drawing Lewis Structures
• Count the total number of valence electrons– An atom’s number of valence electrons is equal to its
group number• Determine the total number of shared electrons
electrons needed – valence electrons present• Connect the atoms with single bonds
– A single bond is one shared pair of electrons• Use lone pairs and/or multiple bonds to give
each atom an octet of electrons
Lewis Structure (Single Bonds)
• Draw Lewis Structures for:• H2O
• HCl
• NH3
Lewis Structures (Multiple Bonds)
• CO2
• N2
Ions• Definition: Ions are atoms or groups of
atoms with an electrical charge• Cations: are positively charged, due to
loss of electrons (Metals)• Anions: are negatively charged, due to
gain of electrons (Non-Metals)• Number of electron’s gained or loss is due
to atoms wanting Octet
Examples of Ions
• Na• Ra• Al• Se• O• Cl• F
Ionic Compounds
• Ionic compounds are held together by ionic bonds, or the attraction of oppositely charged ions
• In the solid state, ionic compounds form crystalline lattices– Cations are attracted to
all the neighboring anions, not just one
– Thus, there are no discrete ionic “molecules”
Ball and Stick Model
Transition Metal Cations• Most transition metals form more than 1 cation
+1 only Ag+
+2 only Zn2+, Cd2+
+1 and +2 Hg22+, Cu+
Hg2+, Cu2+
+2 and +3 Cr2+, Fe2+, Co2+
Cr3+, Fe3+, Co3+
+2 and +4 Sn2+, Pb2+
Sn4+, Pb4+
Polyatomic Ions NO3
-
nitrate
SO42-
sulfate
PO43-
phosphate
NO2-
nitrite
SO32-
sulfite
HPO42-
monohydrogenphosphate
CO32-
carbonate
NH4+
ammonium
H2PO4-
dihydrogenphosphate
HCO3-
BicarbonateOr hydrogen carbonate
OH-
hydroxideC2H3O2
-
acetate
Formulas of Ionic Compounds
• The net charge on a formula unit must be zero
S (+) charges = S (-) charges• Since there are no ionic “molecules” the
formula of an ionic compound is the simplest ratio of cation to anion that gives an electrically neutral combination
Al3+ and O2-
Ca2+ and O2-
Writing Ionic Compound Formulas
• Write the formula for each of the following pairs of ions
• Na and Oxygen• Mg and Fluorine• Rb and Iodine
Nomenclature • Rules for naming compounds and molecules• Anions
– Name the element, drop the ending leaving the root and add “ide”• Element – root + “ide”
• Cl• O• N• S• I
Naming Ionic Compounds1. Name the cation by naming the element
• If the cation is a transition metal you need to distinguish the charge using Roman Numerals• Fe2+ is named Iron (II)• Pb4+ is named Lead (IV)
2. Name the anion • Can be an elemental anion or polyatomic
3. Combine them as two words
Naming Ionic Compounds
K2O
Li2CO3
K2SO4
NaHCO3
Cr2O3
Formulas from Names
• What are the formulas of these compounds?
calcium sulfide
iron (III) acetate
Chromium (III) sulfate
Naming Molecular Compounds• Name each element• Indicate how many of each element is
present with a prefix multiplier– Mono =1; di =2; tr i=3; tetra =4; penta =5;
hexa =6; hepta = 7; octa = 8; nona = 9; deca = 10
• Add the suffix “ide” to the last element• The prefix multiplier mono is left off of the
first element in the compound
Naming Molecular Compounds: Examples
• IBr
• NI3
• N2O4
Formulas from Names
• Sulfur dioxide
• Diphosphorous pentoxide
• Carbon tetrachloride
Molecular Compounds: Common Names
• These compounds have common (non-systematic names)– Water (H2O)
• Dihydrogen monoxide– Ammonia (NH3)
• Nitrogen trihydride– Methane (CH4)
• Carbon tetrahydride– Nitrous oxide (N2O)
• Dinitrogen monoxide– Hydrazine (N2H4)
• Dinitrogen tetrahydride
Acids
• Acids are compounds that can donate an hydrogen ion (H+ ion)
• Acids fall into two categories– Binary Acids HX– Oxoacids HXOn
• Polyatomic anions
Binary Acids
• Most binary acids result from dissolving the corresponding molecular compound in water
• Binary acids are named as hydro (stem name of X) ic acid
HCl(g) HCl(aq)
HCN(g) HCN(aq)
Oxoacids
• Oxoacids are named based on the oxoanion
• “Ate” anion => ic acid• “Ite” anion => ous acid
Oxoacids
CO32- (carbonate anion) H2CO3 (carbonic acid)
NO2- (nitrite anion) HNO2 (nitrous acid)
NO3- (nitrate anion) HNO3 (nitric acid)
PO43- (phosphate anion) H3PO4 (phosphoric acid)
SO32- (sulfite anion) H2SO3 (sulfurous acid)
SO42- (sulfate anion) H2SO4 (sulfuric acid)
Polyatomic Anion
Review of What We Know
• We can write formulas• We can name compounds and molecules• We can draw Lewis Structures
– But what do these molecules look like?
VSEPR Theory
• VSEPR: Valence Shell Electron Pair Repulsion
• Like charges repel and want to be as far apart as possible
• Therefore a given combination of electrons will form into a specific shape
VSEPR
1. Draw the Lewis Structure
2. Assign the central atom (A)
3. Determine the number (n) of atoms bonded to (A) designate them (Xn)
4. Determine the number of lone pairs on (A) designate them (Em)
5. Put together the AXnEm notation
X + E = 2
X + E = 3
X + E = 4
X + E = 5
X + E = 6
VSEPR Examples
• What is the geometry of • CO2
• BF3
• H2O
• NH3
Electronegativity• Linus Pauling developed the
electronegativity scale• Electronegativity is a measure of an
atom’s affinity for electrons• Fluorine is the most electronegative
element (EN=4.0)• The closer an atom is to fluorine,
the more electronegative it is
Polar Covalent Bonds• If two atoms of identical electronegativity are
bonded together, the bond is non-polar• If two atoms of different electronegativity are
bonded together, the bond is polar, and the electrons spend more time around the more electronegative atom– This creates partial charges
• The greater the difference in EN between two atoms, the more polar the bond– The limiting example of this is the ionic bond
EN Type of Bonding0.0 Pure covalent bond
(equal sharing of e- ‘s)0.1 – 0.4 Non-polar covalent bond
(almost equal attraction for shared e- pairs)
0.5 – 1.4 Polar covalent bond (unequal sharing of e- ’s)
1.5 – 3.2 Ionic bond (e- transfer)
Example
• The bond in hydrogen is
• The bond in hydrogen chloride is
Molecular Polarity
• Bond dipoles are vectors• The vectoral sum of the bond dipoles gives the
molecular dipole• Based on the shape of the molecules you can
predict if the dipoles will cancel each other or if they will create a dipole moment
• If a dipole moment exists then the molecule is said to be polar
• If no dipole moment exists then the molecule is said to be non-polar
Molecular Polarity Examples
• Is carbon dioxide polar or non-polar?
• Is water polar or non-polar?
• Is boron trifluoride polar or non-polar?
Intermolecular Forces
• These are attractive forces between molecules or atoms or ions
• Immensely important– These forces hold DNA molecules in a helix
and and are the mechanism for DNA transcription
Dipole Dipole Attraction
• This is the attraction between the opposite (partial) charges of polar molecules
H Cld+ d-
H Cld+ d-
Hydrogen Bonding• This is generally stronger than dipolar
attractions• Hydrogen bonding occurs between a
hydrogen atom and O, N or F.• For H-bonding to happen the H must be
directly bonded to a O, N or F.
O—H
H
O—H
H This is an attraction not really a bond
• Also called Van der Waal’s forces, these are created by instantaneous dipoles
• London forces are much weaker than either dipole-dipole or H-bonding
• London forces get stronger with larger atoms/molecules
London Forces
HeHe
London Forces Between Helium Atoms
HeHe-d +d HeHe-d +d -d +dHeHe+d +d-d -d
HeHe
At a given instant, the electrons on an atommay be non-symmetrically distributed.This leads to creation of a temporary dipole.
As the electrons re-distribute, the dipoles andand the attraction vanishes.
This dipole induces temporary dipoles on neighboring atoms.
For the merest fraction of time, there is adipole-dipole attraction between the atoms.
Ion Dipole Attraction
• This is the attraction between an ionic charge and a polar molecule
• This attraction allows ionic solids to dissolve in water
• The strength of this force varies widely and depends on the magnitude of the dipole moment of the polar species and the size of the ion
A Sodium Ion and a Chloride Ion Hydrated by Water Molecules
Effects of Intermolecular Forces
• More intermolecular forces mean:– Higher boiling and melting points– More viscous liquids
• IM Forces also affect solubility– ‘like dissolves like’
Predicting Boiling Points based on IMF’s
SnH4, CH4, GeH4, SiH4 HBr, HI, HCl, HF
Trends in Boiling Point
H2O
H2S
H2Se
H2Te
Example
• Is carbon dioxide soluble in water? Explain
Example
• Are ionic compounds more soluble in water or in gasoline (a non-polar solvent)? Explain