chemical bonding unit 4 chapters 15 & 16 ionic bonding the bond in ionic compounds (two ions) ...
TRANSCRIPT
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Chemical Bonding
UNIT 4
Chapters 15 & 16
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Ionic Bonding The bond in ionic compounds
(two ions) Held together tightly High melting points
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Compounds are formed from chemically bound atoms or ions
Substances become more stable through chemical bonding, where
2 or more atoms are joined together by a simultaneous
attraction.
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Valence electrons are electrons in the highest occupied energy level of an atom ( the last shell).
Bonding involves only electrons
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Bonding involves only the valenceelectrons.
Na Cl
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i. Ionic Bonding
ii. Covalent Bonding
iii.Metallic Bonding
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Ionic Bonds occur when the more electronegative element “steals” the electron pair away from the other atom.
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The atom that has stolen the electron pair becomes a negativeion (anion) while the “victim” becomes a positive ion (cation).
The two atoms are held together by their opposite charges.
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Can you predict which atoms will gain electrons and which will looseelectrons by looking at the trend in electronegativity?
Increase in Electronegativity
Incr
ease
in
Ele
ctro
neg
ati
vit
y
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When you consider that for an ionic bond to form there must be agreat deal of difference in electronegativity between the atoms, can you predict what two types of atoms allow this to occur?
Metals Non-Metals
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Ionic Properties
Why do most ionic compounds have similar properties?
We can hypothesis that it is due to the bonds formed between the
ions, holding them firmly in a rigid structure
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Forming ions Na and Cl which one will lose
electrons which one will gain electrons
Write out Tin’s electron configuration what will it do??
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Lattice energy The change in energy that
takes place when separated gaseous ions are packed together to form an ionic solid
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The anions and cations in an ionic compound are locked in a regular neutrally charged structure, held by the balance of attractive bonds and electrical repulsion.
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The component ions in such crystals are arranged in repeating three-dimensional (3-D) patterns.
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Alkali metals combine with halogens in 1:1 ratios since alkali metals need
to lose 1 e1- and halogens need to gain 1e1-.
Alkaline earth metals combine with halogens in 1:2 ratios since alkaline earth metals need to lose 2 e1- and
halogens need to gain 1e1-.
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Lewis Structures can be used to illustrate the formation of ionic
bonds.
Be 2 F+ [Be]2+F F1- 1-
Write an equation with electron dot diagrams to illustrate the
formation of aluminum chloride.
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Al 3 Cl+ [Al]3+Cl Cl1- 1-
Cl1-
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Lewis Structures Duet Rule = applies to H and
He and states these two atoms are stable with 2 electrons in their outer shell
Octet Rule= elements are most stable with 8 electrons in their outer shell
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1. Most are crystalline in structure
2. High melting/boiling points
3. Electrically neutral
4. Can conduct electricity when melted or in aqueous solution
5. Hard/ Brittle
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LEWIS DOT STRUCTURES: Elements
Board Practice
Elements #1-20
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Covalent Bonding
Br + Br Br Br
O + O O O
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Covalent Bonding Electrons are shared by
nuclei Polar covalent bonds –
unequal sharing of electrons
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Types of Bonds
1) Single bond – 1 pair of e- are shared- lowest in energy-longest bond length
2) Double bond – 2 pairs of e- are shared
3) Triple bond- 3 pairs of e- are shared
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The most common chemical bond results when the nuclei of 2
atoms are attracted to a pair of shared electrons. If the sharing is equal, because the atoms are
the same, this is called COVALENT BONDING
H HElectron pair
H H
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One atom becomes slightly positive the other slightly negative
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The force of attraction of an element’s nucleus for electrons is called electronegativity (En). Atoms of different elements have different electronegativities. The higher the En, the stronger the attraction for electron pairs.
HFDifference
in En?
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The bonding electrons are on the average closer to the fluorine than to the hydrogen atom.
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The movement of the negatively charged electrons away from hydrogen toward fluorine, due to a difference in electronegativity, builds up a partial negative charge on the fluorine and a partial positive charge on the hydrogen.
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This is not a complete transfer of an electron from hydrogen to fluorine; it is merely a drifting of electrons toward fluorine.
H F
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H FWhen a charge separation of this
type is present, the molecule possesses an electric dipole, and
the bond is called a POLARCOVALENT BOND , or simply
a POLAR BOND.
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H FPolar covalent bond (polar bond) covalent bond joins two atoms of different elements and the bonding electrons are shared unequally
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Resonance
= occurs when more than one valid Lewis structure can be written for a particular molecule
Ex. NO3-1
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Exceptions to the Octet Rule
1) B and Be usually have less than 8 electrons
2) Elements in the 3rd energy level and above can have more than 8 electrons in their outer shell
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Non-polar covalent bond bonding electrons are shared
equally
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1. Soft and squishy
2. Low boiling/melting points
3. Tend to be more flammable
4. Do not conduct electricity
5. Usually non-soluble in water
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Electrostatic attraction force between the cation and free electrons.
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Any successful bonding model for metals must account for the typical physical properties of metals: malleability, ductility, and efficient and uniform conduction of heat and electricity in all directions.
Most metals are durable and have high melting points.
These facts indicate that the bonding in most metals are strong and nondirectional.
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Metal atoms are arranged in very compact and orderly patterns.
i)Body-centered cubic
ii)Face-centered cubic
iii)Hexagonal close-packed
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1. Can conduct electricity (free electrons)
2. Malleable (put into shape)
3. Ductile ( made into wires)
4. Good conductors of heat and electricity
5. Metals are usually shiny
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Lewis Structures for
Molecular Compounds
N N
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Bonding capacity is the number of covalent bonds (shared electron
pairs) that an atom can form.
Covalent molecules often consist of atoms of different elements, with
different bonding capacities.
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How do we decide on their structural arrangement , when we draw structural formulas?
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STEP 1:place the single atom in the center and other atoms around it evenly spaced
CHH
HH
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STEP 2: count the total # of valence e- for all atoms involved in the bonding
Carbon: 1 carbon with 4 valence electrons (1x4) = 4
Hydrogen: 4 hydrogen with 1 valence electrons (4x1) = 4
CH4
4+4 =8
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STEP 3: place the electrons in pairs between the central atom and each non-central atom
C HHH
H
CH4
4+4 =8
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STEP 4: place the remaining electrons around the non- central atom until each has 8 electrons (H atoms have) only 2e-
Ex: AsBr3
AsBr
Br
Br5 + (7x3) = 26
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STEP 5: if electrons remain they are placed in pairs around the central atom
AsBr
Br
Br
Ex: AsBr35 + (7x3) = 26
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STEP 6: if the central atom is in group 14, 15, 16, 17 or 18, the octet rule must be satisfied by moving electron pairs from non-central atoms, creating multiple bonds.Ex: SO2
6 +(6x2)
=18S OO
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1.Central B(6e-) and Be(4e-) will have less than 8 electrons2. If the central atom is in energy level 3 or more it may have more than 8 electrons around it (these energy levels can have 18e-)
S ClCl
ClClSCl4
6+(7x4)
= 34
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NOTE: central atom(s) tend to have the highest bondingcapacities or/and the lowest En.
Draw Lewis Structures for the following:
H2O, NF3, Cl2, SnCl2, PCl5, SO3,
BeCl2, C2H6, C2H2, ClF3, CHCl3, ICl,
O2, N2, SF6, CO2, BF3, C2H4, O3, IF7
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Compounds are arranged in many different shapes
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The VSEPR Theory states that because electron pairs repel, molecular shape adjusts so the valence-electron pairs are as far apart as possible.
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VSEPR model Valence shell electron pair
repulsion Used to predict the geometry
of molecules The structure will minimize
electron pair repulsions
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LINEARLINEAR: the two bonding pairs arrange themselves at 180°. They are connected in a straight line. Ex. CO2, BeF2, HCN, CS2
Groups = 2
Pairs = 0
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A molecule that has a tetrahedral shape has all four pairs of electrons bonded
Group = 4
Pair = 0
The four electron pairs repel each other forming an angle of 109.5°
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Molecules with a bent shape have four pairs of electrons, but only two pairs are bonding pairs (two are lone pairs). Ex H2O, SO2
The bond angle is 109.5 °
Group = 2
Pairs = 2
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A molecule with trigonal planar shape has three bonds all of which lie in the same planeEx. Boron trifluorideThe bond angle are 120 °.
Group = 3Pairs = 0
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A molecule with trigonal pyramidal shape has four pairs of electrons all repelling each other.
Groups = 3Pairs =1
Ex. ammonia
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Intermolecular forces play a key role in determining the physical and chemical properties of covalent compounds.
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Van Der Waals consists of 2 possible types of forces:
1.London Dispersion Forces
2.Dipole-Dipole Forces
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-This is the only type of force present in non-polar covalent molecules.
-It is the weakest of the intermolecular interactions caused by the motion of the electrons.
-The strength of dispersion forces generally increases as the number of electrons in a molecules increases.
Ex. Halogen diatomic molecules.
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- This occurs when polar covalent bonds are attracted to one another.
- Electrostatic attractions occur between oppositely charged regions. (partially (–) and partially (+)).
- Dipole interactions are similar to but much weaker than ionic bonds.
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Dipolar molecules Have a center of positive
charge and a center of negative charge.
aka: dipole moment
Ex. HF
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Dipole moment in NH3
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Dipole cancels out in CO2
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This is found in polar covalent molecules that have hydrogen that is bonded to a very electrostatic element (N2, F2, O2)
Hydrogen bonds are the strongest of the intermolecular forces.
Hydrogen > dipole-dipole > London Dispersion
Bonds interactions forces
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Hydrogen bonds are extremely important in determining the properties of water and biological molecules such as proteins.
The water molecule has a bent shape (105°) and is considered to be polar and the universal solvent.
The attraction in water results from the intermolecular hydrogen bonds.
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Surface tension: the inward force, or pull that tends to minimize the surface area of a liquid
- this surface tension tends to hold a drop of liquid in a spherical shape
The higher the surface tension, the more nearly spherical is the drop of that particular.
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Because of hydrogen bonding, water absorbs a large amount of heat as it evaporates or vaporizes.
The hydrogen bonds must be broken before water changes from the liquid to vapor state.
Vapor Pressure the force exerted due to the gas above the liquid
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Boiling Point: occurs when the temperature at which the vapor pressure of the liquid is just equal to the external pressure.
Boiling leads to evaporation of a liquid. In the case of water, hydrogen bonds break in order for the liquid to vaporize.