Chemistry 1

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• Periodic Table Trends+ /25
• Electron Arrangement
• Atoms: Building Blocks
+ /15
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Next Exam is _____________ !!!
• Chemical Bonding
• Periodic Table Trends
• Electron Arrangement

• Chemical bond is the mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together.
• Why do atoms bond?
• Most atoms are more stable when combined
• When atoms combine, their potential energy decreases and creates more stable arrangements

• All atoms strive to fill their highest energy level with 8 electrons (sometimes 2 electrons).
• They achieve this filled highest energy level by sharing or transferring (gaining/losing) electrons.
• The result is Noble Gas electron configuration
• Exceptions
• Hydrogen (H) wants 2 electrons
• Boron (B) wants 6 electrons (3 pairs)
• Expanded bonding including d electrons

To determine if an atom is going to lose or gain an electron, use the valence electrons to determine which number is easier to get to: 0 or 8.
• Examples:
Gallium:
Iodine:
Cesium:

• Form Ionic Compounds
• Bonding between two oppositely charged ions (cation and anion) that attract when one atom gains/loses electrons to another atom.
• Ions involved can be formed from a single atom (monatomic) or from a molecule (polyatomic)
• High melting/boiling points
• Metal and Nonmetal(s)
• Conduct electricity
• Soluble (usually) in water

Crystal Lattice Structure
• Ionic compounds are typically crystalline solids that form a crystal lattice structure of many ions, NOT individual molecules.
• This structure minimizes potential energy.
Described by a Formula Unit
• In ionic compounds, a formula unit is the simplest collection of atoms from which an ionic compound’s formula can be established.
**Orbital Notation Description of CaCl
2

Form Molecular Compounds (Molecules)
The chemical bond that results when two or more atoms share electrons.
-Positive nuclei of all atoms in bond are sharing the negative electrons
-Covalent compounds are the most common compounds.
-Nonmetal + Nonmetal
- Low melting/boiling points
- Poor conductors of electricity
- Typically insoluble in water

Described by a Molecular Formula
• Indicates the types and numbers of atoms combined in a single molecule of a molecular compound (CO
2
)
• Diatomic molecules are molecules containing only two atoms
(O
2
– Dr. HOFINBrCl)
Two Types of Covalent Bonds
• Non-polar (O
2
)
• Equal sharing of electrons (diatomic elements)
• Electronegativity the same or within 0.3
• Polar (CO)
• Unequal sharing of electrons (one atom attracts the electrons more)
• Electronegativity difference 0.3 to 1.7

• Atoms release energy as they change from atoms to molecules (more stable).
• The amount of energy released (the bond energy) is the difference between the energy of the individual atoms and the energy of the bonded atoms.
**Orbital Notation Description of O
2 and NH
3

• TedEd on Bonding

• Continue/Finish reading Chapter 6 pages 197-
207
• Complete Part ___ of Packet!

Identify the type of bond for the following:
(Ionic, Polar Covalent, Non-Polar Covalent)
• MgCl
2
• CO
2
• Br
2
• BaO
• NH
3

• Covalent bonds usually involve only electrons in outer energy level
• Electron dot notation is used to determine these electrons and visualize how they bond
• Electron dot notation
• Electron configuration notation in which only the valence electrons of an atom are shown using dots around the element’s symbol

1.
Determine the number of valence electrons for the atom.
2.
Write down the symbol of the atom
3.
Begin placing dots around the symbol: start at the top and work around clockwise, placing one dot per side.
4.
Once each side has one dot, you may begin to pair the electrons.
5.
Each side can have no more than two electrons, which means there cannot be more than 8 total valence electrons.

Cs N
Xe Si

Draw the electron dot notation for the following atoms:
1.
Sulfur
2.
Magnesium
3.
Chlorine
4.
Boron

• 1. Place the symbol of the element inside of
brackets.
• 2. Determine the charge of the ion and place the
charge on the outside of the brackets.
• 3. The ions will have either 0 or 8 dots.
* Charge on the ion indicates the number of electrons lost or gained!!

-
1.
Sodium vs. Sodium Ion
2. Beryllium vs Beryllium Ion

• A pair of dots between two atoms represents a bond.
• Other pairs of dots are unshared pairs or lone pairs.
• A dash can represent the two shared electrons
• To demonstrate single, double, or triple covalent bonds use more dashes (pair of shared electrons)

1.
Calculate the total number of valence electrons for the molecule summing the number of valence electrons for each atom.
2.
Write the skeleton structure of the molecule or ion.
Connect every bonded pair of atoms by a dash.
1.
If the number of bonds is higher than the number of places for bonds, then you need double or triple bonds.
2.
Double bonds form between carbon, nitrogen, oxygen, and sulfur.
3.
Triple bonds are typically only seen in carbon and nitrogen.
3.
Distribute the electrons to the atoms surrounding the central atom so each atom has a total of 8 electrons.
(Except Hydrogen)
4.
Distribute the remaining electrons as pairs to the central atom.

1.
HBr
2.
NH
3
3.
SF
2

Q: What is the electron dot formula for COCl
2
?
1. Valence electrons =
2. Skeletal structure =

• Some molecules and ions cannot be represented by a single Lewis structure
• These molecules split time between their different possible structures or resonate
• Together, the structures are called resonance structures
• Bonding in molecules or ions that cannot be correctly represented by a single Lewis structure
• Use a double-headed arrow to illustrate (pg 189 in text)

• Valence Shell Electron Pair Repulsion Theory
• Shape predicted will be close to actual shape
• Basic idea: electron pairs orient themselves so that the repulsions between electron pairs are minimized
• Electron pairs around an atom will be spaced as far apart as possible
• Geometry of Molecules can be described 3 ways: ABE
Geometry, Electron Geometry, and Molecular Geometry

• Description of molecule type, including number of central atoms (A), side/peripheral atoms (B), and lone pairs (E)
• A= Central Atom(s)
• B = Side/Peripheral Atom(s)
• E= Lone Pair(s) on Central Atom

• Geometric arrangement of the shared and unshared electron pairs surrounding the central atom.
• All electron pairs are viewed the “same”.

• The actual 3-D arrangement of a molecule’s atom in space taking into account both shared and unshared electrons.
• Handout

Types of Molecules
• Two atoms (O
2
)
• ABE = AB
• EG = Linear
• MG = Linear (straight line)
• One central element and 2 other atoms, no lone pairs (CO
2
)
• ABE = AB
2
• EG = Linear
• MG = Linear (straight line)
• Three atoms around the central atom, no lone pairs (BF
3
)
• ABE = AB
3
• EG = Trigonal planar
• MG = Trigonal planar (a planar, flat, molecules with 3 atoms at the
• corners of a triangle and a central atom in the middle)

• 4 atoms bonded to a central atom, no lone pairs (CH
4
)
• ABE = AB
4
• EG = Tetrahedral
• MG: Tetrahedral (four sided, keeps them as far apart as possible)
• 5 atoms bonded to a central atom, no lone pairs
• ABE = AB
5
• EG = Trigonal-bipyramidal
• MG = Trigonal-bipyramidal
• 6 atoms bonded to a central atom, no lone pairs
• ABE = AB
6
• EG = Octahedral
• MG = Octahedral

• 2 atoms bonded to a central atom, one lone pair (ONF)
• ABE = AB
2
E
• EG = Trigonal Planar
• MG = Angular/Bent
• 3 atoms bonded to a central atom, one lone pair (NH
3
)
• ABE = AB
3
E
• EG = Tetrahedral
• MG = Trigonal pyramidal molecule (A 3-sided pyramid with triangular
• faces)
• 2 atoms bonded to a central atom, 2 lone pairs (H
2
O)
• ABE = AB
2
E
2
• EG = Tetrahedral
• MG = Angular/Bent

• Packet Practice on Geometries
• Continue reading of Ch. 6 pages 197 to 207

• Outer electrons move around and form a surrounding sea of electrons
• Metal atoms are packed together in crystal lattice structures
• Metal Characteristics:
• – High electrical conductivity and thermal conductivity due to the sea of negatively charged electrons
• – Shiny appearance because electrons are excited to higher levels by absorption of light and then drop back down
• – Malleable and Ductile because one plane of atoms in a metal can slide past another

• Dispersion Forces = the positive end of one momentary dipole is attracted to the negative end of another momentary dipole.
• Dipole = results from uneven charge distribution
• All molecular compounds, even nonpolar
• Dipole-Dipole Forces = attraction between the positive end of one dipole and the negative end of another dipole
• Exist only in polar molecules – increases boiling point
• Only 1/100 as strong as ionic or covalent bonds
• Hydrogen Bonds = Very strong dipole-dipole force
• Occurs between molecules with N-H, O-H, F-H bonds
• Huge difference in electronegativity and small atoms = stronger dipoles
• 10 times stronger than ordinary dipole-dipole forces, still much weaker than covalent or ionic bonds

• Packet Practice on Geometries
• Continue reading of Ch. 6 pages 197 to 207
• Test ______________!!

• Review for test!
• Access notes on blog and in notebook
• Review packet!
• Review Formative Assessments
• Complete ALL remaining parts of the practice packet!