Atoms and Elements: Study Guide for Introduction to Chemistry (CHM1025, Chapter 4)

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Atoms and Elements

Core Definitions

Matter is composed of atoms, and the properties of atoms determine the properties of matter. Understanding the structure and classification of atoms is fundamental to chemistry.

Atom: The smallest identifiable unit of an element.
Element: A substance that cannot be broken down into simpler substances. Each element is defined by its number of protons.
Example: If every atom in a pebble were enlarged to the size of the pebble, the pebble would be larger than Mount Everest.

Additional info: An element consists of one type of atom, but a sample contains many atoms of that element.

Atomic Theory and Its History

Ancient and Modern Theories

The concept of atoms has evolved from philosophical ideas to scientific theory.

Democritus (Ancient Greece): Proposed that matter is composed of indivisible particles called atomos.
Dalton’s Atomic Theory (1808):
1. Each element is composed of tiny, indestructible particles called atoms.
2. All atoms of a given element have the same mass and properties that distinguish them from atoms of other elements.
3. Atoms combine in simple, whole-number ratios to form compounds.
Note: "Matter is neither created nor destroyed" is a conservation principle, not part of Dalton’s three postulates.

Exam Tip: Know Dalton’s three statements and be able to distinguish them from other principles.

The Nuclear Atom and Subatomic Particles

Discovery and Structure

Atoms are composed of smaller particles: protons, neutrons, and electrons.

J. J. Thomson: Discovered the electron, a negatively charged particle much smaller than an atom. Proposed the "plum pudding" model.
Rutherford’s Nuclear Model:
- Most mass and all positive charge are in the nucleus.
- Electrons occupy the space outside the nucleus.
- In a neutral atom, electrons = protons.

Particle	Charge	Location	Relative Mass
Proton	+1	Nucleus	About 1 amu
Neutron	0	Nucleus	About 1 amu
Electron	−1	Outside nucleus	Almost negligible

Additional info: "Subatomic" means smaller than an atom. The nucleus contains protons and neutrons; electrons are found in the surrounding space.

Atomic Number and the Periodic Table

Element Identification and Organization

The atomic number defines the identity of an element and its position in the periodic table.

Atomic Number (Z): Number of protons in the nucleus.
The periodic table arranges elements by increasing atomic number. Elements in the same group (vertical column) have similar properties.
Examples: Silver (Ag) has atomic number 47; Potassium (K) is symbolized from its Latin name kalium.

Element	Symbol	Latin Name
Potassium	K	kalium
Sodium	Na	natrium
Lead	Pb	plumbum
Mercury	Hg	hydrargyrum
Iron	Fe	ferrum
Silver	Ag	argentum
Tin	Sn	stannum
Copper	Cu	cuprum

Periodic Law: Properties of elements recur periodically when arranged by increasing atomic number (Mendeleev).

Metals, Nonmetals, Metalloids, and Important Groups

Classification and Properties

Elements are classified based on their physical and chemical properties.

Class	Location	Properties	Examples
Metals	Left side	Good conductors, malleable, ductile, shiny, lose electrons	Iron, sodium
Nonmetals	Upper right	Poor conductors, varied states, gain electrons	Oxygen, chlorine
Metalloids	Zigzag boundary	Intermediate properties, semiconductors	Silicon, arsenic

Malleable: Can be pounded into sheets.
Ductile: Can be drawn into wires.
Lustrous: Shiny appearance.

Group/Family	Examples	Properties
Alkali metals (1A)	Lithium, sodium, potassium	Lose 1 electron
Alkaline earth metals (2A)	Magnesium, beryllium	Lose 2 electrons
Halogens (7A)	Fluorine, chlorine	Gain 1 electron (−1 charge)
Noble gases (8A)	Helium, neon	Stable, rarely form ions

Additional info: Elements in the same group have the same number of valence electrons, leading to similar chemical behavior.

Ions: Gaining and Losing Electrons

Formation and Charge Calculation

Ions are formed when atoms gain or lose electrons, resulting in a net charge.

Ion: A charged particle.
Cation: Positive ion (loss of electrons).
Anion: Negative ion (gain of electrons).
Ion charge:

Group	Usual Ion Charge	Reason
1A	+1	Lose 1 electron
2A	+2	Lose 2 electrons
3A	+3	Lose 3 electrons
5A	−3	Gain 3 electrons
6A	−2	Gain 2 electrons
7A	−1	Gain 1 electron
8A	0	Stable, do not form ions

Example: Lithium (3 protons, 2 electrons): (Li+); Fluorine (9 protons, 10 electrons): (F−).

Additional info: The nucleus does not change during ion formation; only the electron count changes.

Isotopes and Isotope Symbols

Definition and Notation

Isotopes are atoms of the same element with different numbers of neutrons.

Mass number (A):
Neutrons:
Isotope symbol: AZElement or Element-A (e.g., Ne-20)

Isotope	Protons	Neutrons
Ne-20	10	10
Ne-21	10	11
Ne-22	10	12

Example: Iron-56, Fe3+: Protons = 26, Electrons = 23, Neutrons = 30.

Exam Tip: Isotopes have the same atomic number but different mass numbers.

Average Atomic Mass

Weighted Average Calculation

The atomic mass on the periodic table is a weighted average of all naturally occurring isotopes.

Formula:
Convert percent abundance to decimal (divide by 100).
Multiply each isotope’s mass by its decimal abundance and sum the results.

Example (Gallium):
- Ga-69: 68.9256 amu, 60.11% (0.6011)
- Ga-71: 70.9247 amu, 39.89% (0.3989)
- Calculation: amu

Additional info: The average is weighted toward the most abundant isotope. Do not simply average the masses unless abundances are equal.

Radioactive Isotopes

Definition and Applications

Some isotopes are unstable and emit radiation, making them radioactive. These isotopes have important applications in medicine and science.

Radioactive isotope: An isotope whose nucleus emits subatomic particles (radiation).
Example: Technetium-99 (Tc-99) is used in medical imaging.

Additional info: Radioactive isotopes differ in their rates of decay and biological effects.

Formula Sheet and Exam Checklist

Key Relationships and Skills

Quantity	Rule
Atomic number (Z)	Number of protons
Mass number (A)	Protons + neutrons
Neutrons	A − Z
Neutral atom	Electrons = protons
Ion charge	Protons − electrons
Electrons in cation	Protons − charge magnitude
Electrons in anion	Protons + charge magnitude
Average atomic mass	Σ (fractional abundance × isotope mass)

Be able to define and apply all key terms and relationships above.

Practice Quiz

Sample Questions and Answers

Which subatomic particle has the least mass? Answer: Electron
An atom has 17 protons and 18 electrons. What is its charge, and is it a cation or an anion? Answer: −1; anion
How many protons, neutrons, and electrons are in neutral sodium-23? (Sodium atomic number = 11.) Answer: 11 protons, 12 neutrons, 11 electrons
An ion has 12 protons and 10 electrons. Write its ion symbol. Answer: Mg2+
How many neutrons are in phosphorus-31? (Phosphorus atomic number = 15.) Answer: 16 neutrons
Which family contains chlorine? Answer: Halogens (Group 7A)
Is silicon a metal, nonmetal, or metalloid? Answer: Metalloid
Calculate the average atomic mass for isotopes X-10 (10.0 amu, 20%) and X-11 (11.0 amu, 80%). Answer: 10.8 amu
State Dalton’s rule about how atoms combine to form compounds. Answer: Atoms combine in simple, whole-number ratios.
What is the difference between a mass number and an atomic number? Answer: Atomic number = protons; mass number = protons + neutrons

Summary Table: Key Concepts

Concept	Definition/Rule
Atom	Smallest unit of an element
Element	Substance of one type of atom
Ion	Charged atom (cation or anion)
Isotope	Atoms with same protons, different neutrons
Atomic number (Z)	Number of protons
Mass number (A)	Protons + neutrons
Atomic mass	Weighted average of isotopes