Atoms, Elements, and the Periodic Table: Study Notes for Introduction to Chemistry

Chapter 4: Atoms and Elements

Introduction to Atoms and Elements

Atoms are the fundamental building blocks of matter, and elements are pure substances composed of only one type of atom. Understanding the structure of atoms and the organization of elements is essential in chemistry.

• Atoms: The smallest unit of an element that retains its chemical properties.

• Elements: Substances consisting of only one type of atom, listed in the periodic table.

• Periodic Table: A systematic arrangement of elements based on atomic number and properties.

Historical Development of Atomic Theory

Several scientists contributed to our understanding of atomic structure. Their experiments and models shaped modern chemistry.

• Democritus (460 BC): Proposed that all matter is made up of indivisible particles called atoms.

• Dalton: Developed concepts of atoms, conservation of atoms in chemical reactions, and that atoms combine in whole number ratios to form compounds.

• J. J. Thomson: Discovered electrons using cathode ray tube experiments. Proposed the "Plum Pudding Model" of the atom.

• Ernest Rutherford: Gold foil experiment led to the discovery of the nucleus, showing that atoms have a dense central core.

• Robert Millikan: Oil-drop experiment determined the charge and mass of the electron.

The Atom and Subatomic Particles

Atoms are composed of protons, neutrons, and electrons. These subatomic particles determine the properties of atoms and isotopes.

• Proton: Positively charged particle found in the nucleus.

• Neutron: Neutral particle found in the nucleus.

• Electron: Negatively charged particle found outside the nucleus.

• Atomic Number (Z): Number of protons in an atom; defines the element.

• Mass Number (A): Total number of protons and neutrons in an atom.

• Isotope: Atoms of the same element with different numbers of neutrons.

Definitions and Formulas

• Atomic Number (Z): $Z = \text{number of protons}$

• Mass Number (A): $A = \text{number of protons} + \text{number of neutrons}$

• Isotope: Atoms with the same atomic number but different mass numbers.

• Atomic Mass Unit (amu): Defined relative to 1/12 the mass of a carbon-12 atom.

Example

• Carbon-12: Has 6 protons and 6 neutrons; mass number = 12.

• Carbon-14: Has 6 protons and 8 neutrons; mass number = 14.

Isotopes and Atomic Mass

Isotopes are variants of elements with different neutron numbers. The atomic mass of an element is the weighted average of its isotopes.

• Isotope: Same atomic number, different mass number.

• Atomic Mass: Weighted average of all naturally occurring isotopes.

Example

• Chlorine: Has two main isotopes, Cl-35 and Cl-37. The atomic mass is a weighted average based on their natural abundance.

The Periodic Table

The periodic table organizes elements by increasing atomic number and groups elements with similar properties together.

• Groups (Columns): Elements with similar chemical properties.

• Periods (Rows): Elements with increasing atomic number.

• Metals: Located on the left and center; good conductors, malleable, ductile.

• Nonmetals: Located on the right; poor conductors, brittle if solid.

• Metalloids: Properties intermediate between metals and nonmetals.

Physical Properties Comparison

• Metals: High luster, solid at room temperature, good conductors, ductile, malleable.

• Nonmetals: Gas at room temperature, poor conductors, brittle if solid.

Classification of Elements

• Alkali Metals: Group 1 elements (e.g., Na, K).

• Alkaline Earth Metals: Group 2 elements (e.g., Mg, Ca).

• Transition Metals: Groups 3-12 (e.g., Fe, Cu).

• Halogens: Group 17 elements (e.g., Cl, Br).

• Noble Gases: Group 18 elements (e.g., He, Ne).

Periodic Table Example

Group	Element Type	Examples	Properties
1	Alkali Metals	Li, Na, K	Highly reactive, soft, low melting points
2	Alkaline Earth Metals	Mg, Ca	Reactive, harder than alkali metals
3-12	Transition Metals	Fe, Cu, Zn	Variable oxidation states, good conductors
17	Halogens	F, Cl, Br	Very reactive nonmetals
18	Noble Gases	He, Ne, Ar	Inert, colorless gases

Using the Periodic Table to Determine Ion Charges

The periodic table helps predict the charges of ions formed by elements. Elements in the same group often form ions with the same charge.

• Group 1: Form +1 ions (e.g., Na+).

• Group 2: Form +2 ions (e.g., Ca2+).

• Boron and aluminum: +3 ions

• Group 16: Form -2 ions

• Group 17: Form -1 ions (e.g., Cl-).

• Group 18: Generally do not form ions (noble gases).

Example

• Sodium (Na): Group 1, forms Na+ ion.

• Chlorine (Cl): Group 17, forms Cl- ion.