CHEM S2 FINAL STUDY SET

Electromagnetic Spectrum

All of the frequencies or wavelengths of electromagnetic radiation

Electromagnetic Radiation

Form of energy that exhibits wavelike behavior as it travels through space

Radiant Energy

Energy that travels in the form of waves that have both electrical and magnetic properties

Speed of Light

3.00 x 10⁸ m/s

Electromagnetic Radiation Equation

Speed of light = Frequency x Wavelength

Energy of Light equation

E = h x v

Niels Bohr

Proposed that electrons must have enough energy to keep them in constant motion around the nucleus

Model that Bohr proposed

The planetary model

Ground States

Electrons begin in their lowest energy orbitals

Excited States

When energy is added, electrons absorb it and jump to higher energy orbitals

Energy Levels

Regions of space in which electrons can move about the nucleus of an atom

Electron Cloud Model

shows electrons aren't in fixed orbits but exist as a "cloud" of probability around the nucleus

Electron Cloud

The space around the nucleus where the atom's electrons are found

1st Energy Level

A maximum of 2 electrons

2nd Energy Level

A maximum of 8 electrons

3rd energy level

A maximum of 18 electrons

Valence electrons

Electrons in the outermost energy level

Lewis Dot Diagram

Illustrates valence electrons as dots around the chemical symbol of an element

Periodic law

Physical and chemical properties of the elements repeat in a regular pattern when they are arranged in order of increasing atomic number

Groups

Columns of a periodic table

Periods

Rows of a periodic table

Alkali Metals

Group 1

Alkaline Metals

Group 2

Halogens

Group 17

Noble Gases

Group 18

Transition metals

Group 3 through 12

Metalloids

Found on the "stairstep" boundary between metals and nonmetals

What is the exception to the metalloids

Aluminum

Metals

Easily deformed, good conductors, and have loosely held electrons

Nonmetals

Brittle, poor conductors, and have tightly held electrons

Lanthanides and Actinides

Elements with unfilled F orbitals

What are lanthanides used for

Optical devices

Actinides

All radioactive elements

Atomic radius

The distance from the center of the atom to the outer edge

What are atomic radii typically reported in?

Picometers (1 pm = 10⁻¹² m)

Group pattern for size

Increases

Period pattern for size

Decreases

Ions

Atoms that have gained or lost electrons

Cations

Ions that lose electrons have a positive charge that corresponds to the number of electrons lost

Anions

Ions that gain electrons have a negative charge that corresponds to the number of electrons gained

Ionization energy

Energy required to remove an electron from an atom

Group pattern for energy

Decreases

Periodic pattern for energy

Increases

Electronegativity

A property on an element that indicates its ability to attract electrons in a chemical bond

Electronegativity of noble gases

Zero

Chemical Bonds

Forces that hold two atoms together

Covalent bonding

Attraction between the nucleus and electrons

Ionic bonding

Attraction between positive ions and negative ions

What is the reactivity of atoms based on

Amount of valence electrons

Octet Rule

The tendency of atoms to prefer to have eight electrons in the outermost energy level (valence shell)

Ionic Bond

An electrostatic force that holds oppositely charged particles together

Ionic Crystal

When ionic compounds form, positive and negative ions pack into a regular and repeating pattern

Monoatomic ions

Ions formed from a single atom

Oxidation Number

The charge of an atom

Roman Numerals

Used in the name of transition metals to represent charge

Polyatomic Ions

Ions made up of more than one atom

Formula Unit

The simplest ratio to represent an ionic compound

positive and negative charges have to add up to ______

zero

Chemical nomenclature

Rules for naming compounds

Endothermic reactions

Absorb energy

Exothermic reactions

Release energy

Lattice Energy

energy needed to separate ions of an ionic compound

Properties of ionic compounds

strong attraction of ions, high melting and boiling points, significant hardness, nonconductors as solids, and good conductors when dissolved in water

Delocalized electrons

the electrons involved in metallic bonding that are able to move freely

metallic bond

the attraction of a metallic cation for delocalized electrons

malleable

easily shaped, hammered into sheets

ductile

easily drawn into wire

alloy

a mixture of elements that have metallic properties

Covalent Bonds

A chemical bond is formed from the sharing of valence electrons between two atoms

Molecule

Forms when two or more atoms bond covalently

Diatomic Molecules

Atoms that pair up in nature rather than exist as a single atom

All diatomic molecules

H2, N2, O2, F2, Cl2, Br2, I2

Bonding pair

Shared electrons in a covalent bond

Prefix for 1

Mono -

Prefix for 2

di -

Prefix for 3

Tri-

Prefix for 4

Tetra-

Prefix for 5

Penta-

Prefix for 6

Hexa-

Prefix for 7

Hepta -

Prefix for 8

Octa -

Prefix for 9

Nona-

Prefix for 10

Deca-

Sigma Bond

Single covalent bonds

Pi bonds

Extra bonds formed

Electron Deficient

When there is an insufficient number of electrons to complete the octet of the central atom

Expanded Octet

When an atom can take more than 8 electrons around it

What is the general relationship between bond length and bond strength?

In general, a shorter bond length corresponds to a higher bond strength.

What is the trend for bond length among halogen diatomic molecules (F2, Cl2, Br2, I2)?

F2 < Cl2 < Br2 < I2

What is the trend for bond strength among halogen diatomic molecules (F2, Cl2, Br2, I2)?

F2 > Cl2 > Br2 > I2

How does bond multiplicity (single vs. double vs. triple) affect bond length?

Multiple bonds are shorter than single bonds; triple bonds are shorter than double bonds, which are shorter than single bonds.

Compare the bond lengths of N2, O2, and F2.

N2 (triple) < O2 (double) < F2 (single).

Compare the bond strengths of N2, O2, and F2.

N2 = O2 > F2 (Note: N2 and O2 are both very strong compared to F2).

What is defined as bond dissociation energy?

The energy required to break a specific covalent bond.

Define electronegativity.

The tendency of an atom to attract electrons in a chemical bond.

What determines whether a covalent bond is polar?

The difference in electronegativity between the two bonded atoms.

What is the result of unequal electron sharing in a bond?

An unequal distribution of charge, known as a dipole.

What two conditions must be met for a molecule to be polar?

There must be a difference in electronegativities between atoms, and the molecular arrangement must be unsymmetrical.

Why are diatomic molecules (like F2 or O2) nonpolar?

Because the atoms have identical or similar electronegativities, resulting in equal sharing of electrons.

Can a molecule have polar bonds but still be nonpolar overall?

Yes, if the molecule is symmetrical, the polar bonds cancel each other out (e.g., CCl4).