Chemisty 1B: Midterm 1

H+

Hydrogen Ion

Na+

Sodium Ion

K+

Potassium Ion

Ag+

Silver Ion

Mg²+

Magnesium Ion

Ca²+

Calcium Ion

Zn²+

Zinc Ion

Al³+

Aluminum Ion

Fe²+

Iron (ll) Ion

Fe³+

Iron (lll) Ion

Cu^+

Copper (l) Ion

Cu²+

Copper (ll) Ion

Pb²+

Lead (ll) Ion

Hg²+

Mercury (ll) Ion

NH4+

Ammonium Ion

F^-

Flouride

Cl^-

Chloride

Br^-

Bromide

I^-

Iodide

O²-

Oxide

S²-

Sulfide

N³-

Nitride

NO3^-

nitrate

NO2^-

Nitrite

SO4²-

Sulfate

PO4³-

Phosphate

CO3²-

Carbonate

ClO4^-

Perchlorate

OH-

Hydroxide

CN-

Cyanide

C2H3O2^- or CH3CO2^-

Acetate

Lattice Energy

Strength between 2 charged Particles

2 Factors of Lattice

• Ionic Charge

• Ionic Size

Ionic (blank) makes a greater Contribution to lattice Energy

Ionic CHARGE makes a greater Contribution to lattice Energy

Ionic Charge

Charges of Elements: Direct Relationship

Trend:

• More Charge → More released Energy

• Less Charge → Less Released Energy

Ionic Size

Distance b/w nucleus of atoms: Inverse Relationship

Trend:

• More Distance → Less Released Energy

• Less Distance → More Released Energy

Atomic Radii Trend

Increases going Left and Down

Ionic Bonds are between what elements?

Non-metal x Metal

Ionic Compound Properties

• Solids at Room temp

• High Boiling and Melting Points

• Do not Bend or dent easily

• Conduct Electricity when in water or heated

Covalent Bonds are between what elements?

Non-metal x Non-metal

Covalent Bond Properties

• Exist as solids, liquids, gases: at room temp

• Poor electricity conductors

• Low melting and boiling points

• Energy is RELEASED

Metallic Bonds are between what elements?

Metals

Metal Properties

• Good Electrical Conductivity:

• Are Malleable and Ductile:

• High Melting and Boiling Points

• Luster and Reflectivity: Metals have a shiny appearance (luster) because their free electrons reflect light effectively. This gives metals their characteristic metallic sheen.

• Thermal Conductivity: Metals are good at conducting heat because of the free-moving electrons that transfer thermal energy efficiently through the material.

Intermolecular Forces

Occur Between covalent molecules: weak attractive forces

Intramolecular Forces

Occur Within covalent molecules

3 Kinds of Intermolecular Forces

• London Dispersion

• Dipole - Dipole Moments

• Hydrogen Bonding

Ranking of Forces

Strongest

• Hydrogen Bonds

• Dipole Dipole

• Dispersion

When can Dispersion Forces be stronger?

When a molecule weighs more or has more mass compared to others.

What marks enough difference to determine if Dispersion will be stronger?

0-10: No difference

10-90: maybe

100 < : Dispersion Forces will play big part.

London Dispersion Forces

Weak electrostatic attractions that exist between non-polar covalent molecules or monatomic gases

Temporary Dipole

Dipole-Dipole Moment

a molecule having a slightly positive end and a slightly negative end, caused by uneven sharing of electrons in a covalent bond

Polar Ranges

Periods 1-2

• 0-0.4: Non-polar

• 0.5 - 2: Polar

• 2 < : Ionic

Periods 3 and up

• 0: Non-Polar

• 0.1 - 2: Polar

• 2 < : Ionic

Hydrogen Bonding

Covalent molecules that have a hydrogen atom directly bonded to a highly electronegative atom, such as

• Fluorine, Oxgyen, Nitrogen

Bond Acceptor

Neighboring molecules Oxygen, Nitrogen, Fluorine atoms

Bond Donor

Hydrogen Atom

The (blank) atoms in H(black) molecules act as hydrogen bond acceptors

to neighboring H(blank) molecules, but usually not to any other molecules because of immiscibility or alternative reactivity.

The F atoms in HF molecules act as hydrogen bond acceptors
to neighboring HF molecules, but usually not to any other molecules because of immiscibility or alternative reactivity.

Crystalline Solid

molecules, ions, or atoms have a highly ordered three-dimensional geometric arrangement that is distinctive for that solid.

• Highly Ordered

• Characteristic geometric shapes

• Flat surfaces (or faces) that make distinctive angles with one another

Amorphous Solids

molecules, ions, or atoms are arranged in an irregular manner

• lack the characteristic long-range order found in crystals

• do not display well-defined shapes or flat faces

Crystalline solids have (blank)
melting temperatures

Amorphous solids, on the other hand, do (blank) have well-defined melting temperatures

Crystalline solids have specific
melting temperatures.

Amorphous solids, on the other hand, do not have well-defined melting temperatures.

Polycrystalline Solids

Crystals fused together

Molecular Solids

Covalent that contain both intramolecular forces and intermolecular forces

Discrete Molecules

Molecular Solid Force Strength

Relatively Weak soft substances with low melting points

Ionic Solids

Ionic Compounds, contain extended 3 dimensional array of oppositely charged ions. (cation and anions)

• Held by strong electrostatic attraction

Ionic Solid Force Strength

Hard and Brittle

High Melting Points

Atomic Solids

Bonding between

• Metallic Solids

• Network Solids

• Group 8a solids (noble gases)

Metallic Solids

Metals held together my metallic bonding

• Electrons are delocalized, mobile and free to move throughout solid

• Good conductors of heat and electricity

Network Solids

Large number of non-metal atoms held together by network of covalent bonds.

Group 8A Solids

Noble Gas elements form solid weak London dispersion forces bind to atoms together at low temperatures

Cubic Array Arrangment Cubes

• Simple Cubic Structure

• Body Centered Cubic Structure

Simple Cubic Structure (SC)

Atoms: 1

Coordination Number: 6

52% volume occupied

Equation: 2r

Body Centered Cubic Structure (BCC)

Atoms: 2

Coordination Number: 8

68% Volume Occupied

Equation: 4r/√3

Hexagonal Array Shapes

• Face Centered Cubic (FCC) & Cubic Close Packed (CCP)

• Hexagonal Close Packed (HCP)

Face centered Cubic (FCC) & Cubic Close Packed (CCP)

• Atoms: 4

• Coordination #: 12

• 78% volume occupied

• Equation: l = √8r

Hexagonal Close Packed (HCP)

• Atoms: 2

• Coordination #: 2 atoms

• 78% volume occupied