chemical bonding · number of outer electrons softer metals can be combined with harder metals....
TRANSCRIPT
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CHEMICAL BONDING Unit 4
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BONDING
Chemical Bond: mutual electrical attraction between nuclei and valence electrons of different atoms
Type of bond depends on electron configuration and electronegativity
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WHY DO ATOMS BOND?
TO BECOME STABLE!!Key: 8 valence electrons
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TYPES OF BONDINGIonic: an electrostatic force of attraction between positive and negative ions; ions for when atoms lose/gain electrons (electrons are transferred from metal to nonmetal)
Covalent: sharing of electron pairs between atomspolar unequal sharing of electronsnonpolar equal sharing of electrons
Metallic: an attraction between metal atoms and outer mobile electrons
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COVALENT BONDING
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COVALENT BONDING
Occurs between NONMETALS
Nonmetals have high electronegativities and want to gain more electrons. They cannot lose valence electrons. In order to bond, nonmetals must then SHARE valence electrons.
The atoms share enough electrons to obtain 8 valence electrons (including the shared electrons)
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TYPES OF COVALENT BONDS
Equal sharing of electron between atoms in a compound, no positive or negative poles exist: NONPOLAR!
Unequal sharing of electrons, positive and negative ends exist: POLAR!
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DETERMINE BOND FROM ELECTRONEGATIVITY DIFFERENCES
0 0.3 1.7 3.3
Covalent Ionic
Non-polar Polar
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DETERMINE THE TYPE OF BONDUse the following electronegativity values to determine the bond character of each:
H – S (2.2, 2.6)
S – Cl (2.6,3.2)
Rb – Se (0.8,2.6)
C – I (2.6,2.7)
Cs – S (0.8, 2.6)
O – O (3.4,3.4)
F – I (4.0,2.7)
Ca – N (1.0,3.0)
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OTHER COVALENT BONDSCovalent network bonding: 3-D network of covalently bonded atoms macromolecules
Very high melting points, very hard solids
Ex: diamond, graphite, quartz
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COORDINATE COVALENT BOND
DEF: When both of the shared electrons in a covalent bond come from the same atom
Coordinate Covalent bond will form here
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TERMS TO KNOW!Molecular compound: neutral compound consisting of nonmetals covalently bonded in which the electrons are shared
Molecule: smallest representative unit of a molecular compound, can exist independently
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DIATOMIC MOLECULES (MEMORIZE THESE!)
Diatomic molecules: molecules consisting of two atoms of the same element. They are ALWAYS found in pairs unless in a compound with something else
Ex: H2 O2 F2 Br2 I2 N2 Cl2“HOFBrINCl”
**Form a “7” starting at atomic number (Z) 7 and include H
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DEFINITIONS
Bond length: average distance between nuclei of two bonded atoms (sum of atomic radii)
Bond angle: angle between two bonds in a molecule
Bond energy: energy needed to break a bond and form neutral atoms
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FACTS
As bond length increases, bond energy decreases.
As bond length increases, bond strength decreases.
**Think about it!!!**
What relationship do each of the above have?
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APPLY THAT INFORMATION!
Question: What can you say about the bond strength and bond length of these 2 compounds?
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POLYATOMIC IONS
DEF: covalently bonded group of atoms that have an overall charge.
Examples: Nitrate (NO3-), Hydroxide (OH-), Ammonium (NH4
+)
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PROPERTIES OF MOLECULES
Low melting & boiling pointsBrittle, dull solids or gasesPoor conductors of heat and electricity
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WARM UP – 2/25
You are given a covalent compound. Answer the following:
Will it conduct electricity
Is it malleable
Could you melt it in your kitchen?
What is it doing with its electrons?
Are there any metallic atoms in it?
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OCTET RULEDefinition: atoms tend to gain, lose or share electrons so that they have eight valence electrons
exceptions Hydrogen (2e-), Beryllium (4e-), Boron (6e-)
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STEPS DRAWING LEWIS STRUCTURES1. Total the number of valence electrons in a molecule.
a) If there is an overall positive charge, subtract it from the number of valence electrons.
b) If there is an overall negative charge, add it to the number of valence electrons.
2. Draw the atom in the molecule connected with a single bond. If carbon is present, it is the center atom. Otherwise the atom in the singlemost form is the center atom.
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STEPS DRAWING LEWIS STRUCTURES3. Add electrons so that all of the atoms are happy. (Octet
Rule)4. Compare the number of electrons you have drawn to the
number of valence electrons.a) If you have drawn too many electrons, make double or
triple bonds. In order to make a multiple bond, you must take away a pair of electrons from the atom on each side of the bond.
b) If you have not drawn enough electrons, add extra electrons to the central atom.
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LET’S TRY THESE EXAMPLES!
NH3
CH4
H2O
BF3
CO2
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MOLECULAR GEOMETRY a.k.a. Molecular Shape
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VSEPR THEORY
Valence Shell Electron Pair Repulsion electron pairs (clouds) spread as far apart as possible to minimize repulsive forces
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IMPORTANT VSEPR TERMS
Lone Pair: A pair of electrons that do NOT create a chemical bond
Shared Pair: A pair of electrons that make a chemical bond between elements Lone Pair of
electrons in an electron cloud
Shared Pair of electrons
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VSEPR Shape # Bonds to Central Atom
# Lone Pairs to Central
Atom
Examples
Linear 2 atoms together or 2
bonds to central atom
0 H-ClO = C = O
Bent 22
21
Trigonal Planar
3 0
Trigonal Pyramidal
3 1
Tetrahedral 4 0
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LET’S TRY IT - PRACTICE PROBLEMS!
NH3
CH4
H2O
BF3
CO2
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POLARITY
Molecular polarity is the distribution of molecular charge (even or uneven)
Molecular polarity depends on:SymmetryMolecular shapeLone pairs
Molecular polarity influences intermolecular forces.
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POLAR VS NON POLAR STRUCTURES
Water is polar
Methane is non polarHydrogen Fluoride is polar
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MOLECULAR POLARITY
Nonpolar molecules are usually symmetrical. (about their bonds)
Polar molecules are usually nonsymmetrical. (about their bonds)
Polar molecules are called dipoles. (positive and negative ends)
**As soon as you have a lone pair on your central atom, your molecule is polar.**
Two shapes are always polar: bent and trigonal pyramidal.
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LET’S TRY IT - PRACTICE PROBLEMS!
NH3
CH4
H2O
BF3
CO2
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INTERMOLECULAR FORCES
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INTERMOLECULAR FORCES
Definition: weak forces of attraction between molecules
TypesDipole-dipole forcesHydrogen bondingLondon Dispersion forces
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INTERMOLECULAR FORCES
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DIPOLE-DIPOLE FORCES
Exists between polar molecules
Results in higher melting point and boiling point than expected
Substances exist mostly as solids or liquids
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HYDROGEN BONDING
DEF: Special dipole-dipole force occurring when hydrogen on one molecule is attracted to N, O, F of another molecule
Ex: gives water its unusual properties, ice floats in liquid water, higher melting/boiling point, surface tension
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HYDROGEN BONDING
http://www.chemguide.co.uk/atoms/bonding/hbond.html
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LONDON DISPERSION
DEF: instantaneous dipole due to a shift in electrons in the molecule
Occurs between Nonpolar-Nonpolar
Strength increases with an increase in number of electrons
Substances with LDF only have low melting points/boiling points and are mostly gases
Ex: diatomic molecules, noble gases
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DETERMINE THE IMF(S) PRESENT
1) H2O2) SCl23) PF3
4) CF4
Helpful Hint: Draw the molecular structure and determine if the bond is polar or non polar.
DispersionForces
Dipole-Dipole
Hydrogen Bonding
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HOW DO THESE PROPERTIES RELATE TO BOND TYPE OR INTERMOLECULAR FORCES?Stronger bonds/imf’s higher melting point/boiling point
Weaker bonds/imf’s lower melting point/boiling point
Solubility “likes dissolves like”Polar dissolves polar, nonpolar dissolves nonpolar
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METALLIC BONDING
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METALLIC BONDING
Definition: an attraction between metal atoms and outer mobile electrons (delocalized electrons)
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METALLIC BONDINGDEF: an attraction of metallic atoms for “delocalized electrons;” electrons roam traveling freely from one atom to another
positive ions in a sea of electrons
http://www.revisioncentre.co.uk/gcse/chemistry/bonding.html
http://www.ndt-ed.org/EducationResources/CommunityCollege/Materials/Structure/metallic.htm
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FUN FACTSThe strength of a metallic bond is determined by number of outer electrons softer metals can be combined with harder metals. (forming alloys)The strongest metals are the transition metals.Some “d” electrons are delocalized and involved in bonding
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PROPERTIES OF METAL CRYSTALS
Good conductors of heat and electricityShiny (luster)Hard, metallic crystalsMalleableDuctileHigh melting point
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WARM UP – 2/28
Tell me if the following compounds are covalent, ionic, or metallic?
CH4 SO3
NaCl FeO2
K2SO4 Ag
Fe NH3
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IONIC BONDING
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IONIC BONDING
Definition: an electrostatic force of attraction between positive and negative ions; ions form when atoms lose/gain electrons
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IONIC BONDINGIonic compounds are formed by the force of attraction between positive and negative ions. It involves a transfer of electrons from metal to nonmetal
# (+) = # (-) (metals – nonmetals)
Formula unit: simplest combining ratio of ions in a compound does not exist independently
http://www.visionlearning.com/library/module_viewer.php?mid=55
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DETERMINING THE FORMULA UNIT
Key point: overall charge must be neutral!Potassium and fluorine
Calcium and iodine
Magnesium and nitrogen
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LET’S TRY IT!
Calcium and nitrate
Sodium and phosphate
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STRUCTURES - DOT DIAGRAM FOR NaCl
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STRUCTURES - DOT DIAGRAM FOR CaBr2
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IONIC COMPOUNDSCrystal lattice: 3-D crystal structure (arrangement of ions)
Lattice energy: energy released when one mole of ionic crystalline compound is formed from gaseous ions
http://www.visionlearning.com/library/module_viewer.php?mid=55
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PROPERTIES OF IONIC COMPOUNDS
High melting point Most soluble in waterCrystalline solidGood conductors in liquid or aqueous stateHard solids but will fracture
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REVIEW ACTIVITY Ionic or covalent?
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LITHIUM BONDING WITH NITROGEN
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TWO NONMETALS BONDING
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FLUORINE BONDING WITH MAGNESIUM
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METAL BONDING WITH NONMETAL
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Cs2S
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PH3
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FBr3
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CALCIUM AND PHOSPHATE
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LOW MELTING POINT
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HF
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METALS ONLY
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SHARING ELECTRONS
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DELOCALIZED ELECTRONS
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CCl4
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H2S
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Na3N
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WARM UP – 3/1
Draw the following structures. If applicable, determine the shape and the strongest IMF present.
SiS2
Mg2N3
H2S
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WARM UP – 3/4
Tell me the three MOST DIFFICULT things from this unit. (I said MOST so you can’t say “everything”)