Atoms, Elements, and Chemical Bonding: Study Guide

Study Guide - Smart Notes

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Elements and the Periodic Table

Pure Substances and Element Symbols

A pure substance is a material made of only one type of particle, either an element or a compound. Elements are represented by unique symbols, often derived from their English or Latin names. Students should know the symbols for Elements 1-20, plus Fe (Iron), Ni (Nickel), Cu (Copper), Zn (Zinc), Br (Bromine), Ag (Silver), Sn (Tin), I (Iodine), Au (Gold), Hg (Mercury), and Pb (Lead).

Element Symbol: One or two letters, first always capitalized.
Example: Carbon is "C", Sodium is "Na" (from Latin Natrium).

Periodic Table Arrangement

The Periodic Table organizes elements by increasing atomic number and groups them by similar properties.

Groups: Vertical columns; elements in a group share similar chemical properties.
Periods: Horizontal rows; elements in a period have the same number of electron shells.
Representative Elements: Groups 1, 2, and 13-18.
Transition Elements: Groups 3-12; often metals with variable charges.
Alkali Metals: Group 1 (except Hydrogen); highly reactive.
Alkaline Earth Metals: Group 2; less reactive than alkali metals.
Halogens: Group 17; very reactive nonmetals.
Noble Gases: Group 18; inert gases.
Metals: Left and center of table; conduct electricity, malleable.
Non-metals: Right side; poor conductors, often gases or brittle solids.
Metalloids: Border between metals and non-metals; properties intermediate.

Elements in a group have similar valence electron configurations, leading to similar chemical behavior. Elements in a period have the same number of electron shells but differ in properties as you move across.

Atomic Structure

Atoms and Subatomic Particles

An atom is the smallest unit of an element, composed of protons, neutrons, and electrons.

Proton: Positively charged particle in nucleus; mass ≈ 1 amu.
Neutron: Neutral particle in nucleus; mass ≈ 1 amu.
Electron: Negatively charged particle; orbits nucleus; mass ≈ 0.0005 amu.

Dalton’s Atomic Theory (early 1800s):

All matter is made of atoms.
Atoms of the same element are identical.
Atoms cannot be created or destroyed in chemical reactions.
Atoms combine in simple ratios to form compounds.

Atomic Mass Unit (amu)

The atomic mass unit (amu) is a standard unit for atomic mass, based on 1/12 the mass of a carbon-12 atom.

Proton: 1 amu
Neutron: 1 amu
Electron: ≈ 0.0005 amu

Structure of an Atom

Atomic Number (Z): Number of protons in nucleus; defines the element.
Mass Number (A): Total number of protons and neutrons.
Isotopes: Atoms of same element with different numbers of neutrons.

Example: Carbon-12 has 6 protons and 6 neutrons; Carbon-14 has 6 protons and 8 neutrons.

Atomic Mass vs. Mass Number

Mass number is the sum of protons and neutrons in a single atom. Atomic mass is the weighted average mass of all isotopes of an element, found on the periodic table.

Nuclear Symbols

Nuclear symbols show the element, mass number, and atomic number. Example:

A: Mass number
Z: Atomic number
X: Element symbol

To find protons, neutrons, and electrons:

Protons = Z
Neutrons = A - Z
Electrons = Z (for neutral atom)

Electron Energy Levels and Periodic Trends

Electron Energy Levels

Electrons occupy energy levels (shells) around the nucleus, described by the principal quantum number (n).

Lowest energy level is closest to nucleus.
Electrons can move between levels by absorbing or releasing energy.
Each level can hold a maximum number of electrons:

Example: First energy level (n=1) holds 2 electrons; second (n=2) holds 8.

Valence Electrons and Group Numbers

Valence electrons are electrons in the outermost energy level. The group number for representative elements indicates the number of valence electrons.

Group 1: 1 valence electron
Group 2: 2 valence electrons
Group 17: 7 valence electrons
Group 18: 8 valence electrons (except Helium, which has 2)

Electron Dot Structures

Electron dot structures (Lewis structures) show valence electrons as dots around the element symbol.

Example: Sodium (Na): one dot; Oxygen (O): six dots.

Periodic Trends: Atomic Size and Ionization Energy

Periodic trends describe how properties change across the table.

Atomic Size: Increases down a group, decreases across a period.
Ionization Energy: Energy required to remove an electron; decreases down a group, increases across a period.

Ionization Energy is the energy needed to remove the outermost electron from an atom in the gas phase.

Chemical Bonding

Ionic and Covalent Bonds

Atoms form bonds to achieve stable electron configurations (usually an octet).

Ionic Bonds: Formed between metals and nonmetals; electrons are transferred.
Covalent Bonds: Formed between nonmetals; electrons are shared.

Example: Sodium chloride (NaCl) is ionic; water (H2O) is covalent.

Valence Electrons and Electron Dot Structures

Valence electrons determine bonding behavior. Electron dot structures can be used for elements and compounds with single bonds.

Octet Rule

Atoms tend to gain, lose, or share electrons to have eight in their outer shell (except for hydrogen and helium).

Ions: Cations and Anions

Cation: Positively charged ion (lost electrons); usually metals.
Anion: Negatively charged ion (gained electrons); usually nonmetals.

Example: Sodium forms Na+; Chlorine forms Cl-.

Ionic Compounds: Writing and Naming

Ionic compounds are formed from cations and anions. The formula is written with the cation first, followed by the anion.

Example: Magnesium chloride: Mg2+ and Cl- combine to form MgCl2.
Transition Metals: May have variable charges; use Roman numerals in names (e.g., Iron(III) chloride).

Characteristics of Ionic Compounds:

High melting and boiling points
Conduct electricity when dissolved in water

Polyatomic Ions

Polyatomic ions are groups of atoms with a charge. Know the formula, name, and charge for:

Hydroxide: OH-
Ammonium: NH4+

Other polyatomic ions may be provided on exams.

Covalent Bonds and Diatomic Compounds

Covalent bonds involve sharing electrons. Diatomic compounds are molecules made of two atoms of the same element.

Diatomic Elements: H2, N2, O2, F2, Cl2, Br2, I2

The number of covalent bonds an element forms is related to its valence electrons.

6 valence electrons: forms 2 bonds
7 valence electrons: forms 1 bond
4 valence electrons: forms 4 bonds

Covalent Compounds: Writing and Naming

Covalent compounds use prefixes to indicate the number of atoms (e.g., carbon dioxide, CO2).

Electronegativity and Bond Polarity

Electronegativity is the ability of an atom to attract electrons in a bond. The difference in electronegativity determines bond type:

Small difference: nonpolar covalent
Moderate difference: polar covalent
Large difference: ionic

To predict bond polarity, calculate the difference in electronegativity values.

Sample Table: Mass Number Relationships

The following table helps determine missing atomic information:

Element	Atomic Number (Z)	Mass Number (A)	Protons	Neutrons	Electrons
Carbon-12	6	12	6	6	6
Carbon-14	6	14	6	8	6
Magnesium-24	12	24	12	12	12
Additional info: If electrons are missing, use atomic number for neutral atoms; if neutrons are missing, subtract atomic number from mass number.

Sample Table: Polyatomic Ions

Ion Name	Formula	Charge
Hydroxide	OH	-1
Ammonium	NH4	+1
Additional info: Other polyatomic ions may include nitrate (NO3-), sulfate (SO42-), etc.