January 31 February 1
Valence electrons through properties of ionic compounds
HW: WS#1
Metallic bonding through Lewis Dots
HW: WS #2
2 3 4 5 6
VSEPR shapes
Class Work: WS#3
Bonus: Chem. Valentines-Due Mon.
Complete notes
Class Work: WS #4
HW: Draw Lewis structures for lab
Covalent Models Lab
Ch. 16 QUIZ
Ch. 15 & 16 Review Activity
Chapter 15/16 Review
Test
Monday!!!
Chemistry I
Chapter 15 & 16: Bonding
15.1 Electron Configurations in Ionic Bonding
Valence Electrons – the electrons in the highest occupied energy level of an element’s atoms
To find the number of _______________________ look at the element’s group number
Electron dot structure - System of arranging dots representing valence electrons. G.N. Lewis developed this system and so they are also referred to as ___________ Dot Structures.
EXAMPLES:
Write the electron configuration for phosphorous.
Write the Lewis dot structure for phosphorous.
Practice
Write Lewis dot structures for the following atoms.
Mg S Cl Rb Si Ar
Octet rule – in forming compounds, elements tend to achieve the electron configuration of a noble gas.
Metals will typically _____________________ and therefore become positive
Na ___ ___ ___ ___ ___ ___
1s 2s 2p 3s
Na+
Write the number of valence
Nonmetals will typically ____________________ and therefore become negative
O ___ ___ ___ ___ ___ ___
1s 2s 2p 3s
O
2-Electron dot symbols, Lewis dot structures, can be used to __________________________. EXAMPLES:
Draw in the dots for each atom in their ion form.
Na
Cl
F
Mg
O
Al
15.2 Ionic Bonds
Ionic Bond - A chemical bond formed by the electrostatic attraction between a cation and an anion.
___________ compounds are generally a metal and a nonmetal bonded together.
EXAMPLE:
When potassium reacts with chlorine what kind of compound is formed? Is there a transfer of electrons?
Using the Lewis structures show what happens to K and Cl when they combine to form salt.
K
+
Cl
K
Cl
Properties of Ionic compounds
High _________________ points
Solids at room temperature
_________________________
15.3 Bonding in Metals
Metallic bonds – consist of the attraction of free-floating valence electrons for the positively charged _______________________________.
These forces hold metals together
Good _____________________________
____________________ – can be drawn into a wire
Malleable – can be hammered into ____________ or forced into shapes
Crystalline structure of metals and ionic compounds
Metals and _____________ compounds are organized in compact orderly patterns Label the
2- Identically sized spheres have several arrangements. Sketch each arrangement in the boxes provided.
1) Simple cubic
2) Body-centered cubic 3) Face-centered cubic
Alloys – mixtures of two of more elements, at least two of which are metals.
________________ are the most important alloys today
Consist of Fe, C, B, Cr, Mn, Mo, Ni, W, V
Chemistry I
Chapter 16: Covalent Bonding
16.1
Covalent Bonds
Covalent bond – Occurs when a pair of __________________ is shared between two atoms.
Often between two nonmetals.
____________ covalent bond – two atoms share a single pair of electrons
Double covalent bond – two atoms share two pairs of electrons
Triple covalent bond – two atoms share three pairs of _____________________
Lewis dot structures can be useful for representing covalent bonds between elements in a covalent compound.
EXAMPLES:
H
+
H
H H
Cl
+
Cl
Cl Cl
O + O O O
N + N N N
Rules for writing electron dot structures: (use pencil!!!)
1. Add up the valence electrons from all the atoms in the compound. Don’t try to keep track of which electrons come from which atoms. If you are working with an ion, you must add or subtract electrons to account for the charge.
2. Put the element that you have the fewest of as the central element. Arrange the remaining atoms around the central atom as symmetrically as possible.
3. Use a pair of electrons to form a bond between each pair of atoms.
4. Arrange the remaining electrons to satisfy the duet rule for hydrogen and the octet rule for all remaining atoms.
What kind of bond would exist between the following?
F2
P2
5. Count the number of electrons represented in the drawn molecule. If two too many electrons are represented, draw a double bond between two elements and remove a pair of electrons form each element taking place in the bond.
EXAMPLES:
H2O CH4 CO2
NH3 CCl4 NH4+
SO42- CN- CO3
2-Exceptions to the octet rule
BF3 PCl5 SF6
_____________________ structures – structures that can occur when it is possible to write two or more valid electron dot structures that satisfy the octet rule.
EXAMPLES:
CO32- NO3
-VSEPR – valence shell ___________________________ repulsion theory
Valence electrons on the central atom repel each other.
Regions of electron ____________________ (where pairs of electrons are found) can be used to determine the shape of the molecule.
EXAMPLES:
CO
2 Why do Here there are two regions of electron density. The regions want to be as far apart as possible, so it is _________________.
CH
4 There are four electron pairs. You would expect that the bond angles would be 90, but because the molecule is three-dimensional, the angles are 109.5. The molecule is of _______________ arrangement.
NH
3 Here there are also four regions of electron density, but because one of the electron pairs is a lone pair, the shape is called ___________________________.
H
2O
There are four regions of electron density, but two are lone pairs. This structure is referred to as _________.
CO
32- There are ____________ regions of electron density. This structure is referred to as trigonal planar.
Draw the Lewis structures below and practice determining molecular shape.
H2S SO2 CCl4
BF3 NF3
16.3 Polar Bonds and Molecules
In covalent bonds, the _______________ of electrons can be equal or it can be unequal.
___________ covalent bond - This is a covalent bond in which the electrons are shared equally.
EXAMPLES:
H2 Br2 O2 N2 Cl2 I2 F2
Polar covalent bond - This is a covalent bond that has a dipole. It usually occurs when 2 different elements form a covalent bond.
A ________________________ has 2 separated equal but opposite charges.
EXAMPLE: HCl
H
+
Cl
H Cl
Electronegativity - This is the measure of the __________________ an atom has for a shared pair of electrons in a bond.
Electronegativity values increase across a period and up a group.
Electronegativity values can be used to determine the degree of electron sharing which demonstrates the type of bond that will occur.
Difference > ____ Ionic Bond
Difference ____ - ______ Polar Covalent Bond
Difference = ____ Nonpolar Covalent EXAMPLE:
Identify the type of bond for each of the following compounds.
HBr NaF N2
Attractions between Molecules
Van der Waals forces – the _____________ of the intermolecular forces. This includes London dispersion forces and dipole-dipole forces.
o London dispersion forces – an attraction between _______________ molecules that is caused by the movement of electrons to form areas of concentration
o Dipole interactions – an attraction between _____________ molecules that is caused by a difference in electronegativity.
Hydrogen bonds – attractive forces in which hydrogen, covalently bonded to a very electronegative atom (N, O, or F only) is also weakly bonded to an unshared (lone) pair of electrons on another ____________________________atom. It may be in the same molecule or in a different, nearby molecule.
Ionic bonding – occurs between metals and nonmetals when electrons are transferred from one atom to another. These bonds are __________________________.
The Strength of Attractive Forces
Ionic bonds > hydrogen bonds > dipole-dipole attractions > LDF Determine if
these are ionic, polar covalent, or nonpolar bonds.
C-O (difference of 1.0)
K-Cl (difference of 2.2)
