Atoms and the Periodic Table

Topic 8: Atoms and the Periodic Table

Overview of Elements

• The periodic table displays the known chemical elements organized in order of increasing atomic number.

• Each entry in the table references the element's atomic number, symbol, and average atomic mass.

Element Listings

Common Elements and Their Symbols:

• C (Carbon) - Atomic Mass: 12.0

• H (Hydrogen) - Atomic Mass: 1.01

• Li (Lithium) - Atomic Mass: 6.94

• Na (Sodium) - Atomic Mass: 22.99

• Mg (Magnesium) - Atomic Mass: 24.30

• N (Nitrogen) — Atomic Mass: 14.01

• O (Oxygen) — Atomic Mass: 15.99

• F (Fluorine) — Atomic Mass: 18.99

• Ne (Neon) — Atomic Mass: 20.18

• P (Phosphorus) — Atomic Mass: 30.97

• S (Sulfur) — Atomic Mass: 32.06

• Cl (Chlorine) — Atomic Mass: 35.45

• Ar (Argon) — Atomic Mass: 39.95

• K (Potassium) — Atomic Mass: 39.10

• Ca (Calcium) — Atomic Mass: 40.08

• Fe (Iron) — Atomic Mass: 55.85

• Au (Gold) — Atomic Mass: 196.97

• U (Uranium) — Atomic Mass: 238.05

Lesson 1: Atomic Theory

Guiding Questions:

• What are the parts that make up an atom?

• What is Atomic Theory?

• What evidence supports the modern model of the atom?

Historical Perspectives on Atomic Theory

• Democritus (430 BCE)

  • Proposed that matter could not be divided indefinitely.

  • Coined the term atomos, meaning "uncuttable".

• John Dalton (1803)

  • Conducted experiments leading to the foundation of modern atomic theory.

  • Core beliefs of Dalton's Atomic Theory:

  • All elements consist of atoms and cannot be divided.

  • Atoms of the same element are identical.

  • An atom of one element cannot be transformed into an atom of another element.

  • Compounds are formed when atoms of different elements combine in specific ratios.

• JJ Thomson (1897)

  • Discovered electrons, negatively charged particles within an atom.

  • Proposed the Plum Pudding Model: a positive core with negative electrons scattered, likened to chocolate chip ice cream.

• Ernest Rutherford (1911)

  • Conducted the Gold Foil Experiment, which involved shooting positively charged particles at gold foil.

  • Found that some particles were deflected, concluding that atoms consist mostly of empty space with a dense positive nucleus (identified as protons).

• Niels Bohr (1913)

  • Suggested a planetary model where electrons move in specific orbits around the nucleus.

  • Each electron exists in a fixed energy level.

• Cloud Model (1920’s)

  • Proposed electrons move in a cloud-like region around the nucleus, not in fixed orbits.

• James Chadwick (1932)

  • Discovered neutrons, another particle found in the nucleus of an atom.

Development of Atomic Theory

Timeline of Contributions:

• 1803: John Dalton - Atoms serve as solid spheres with distinct properties.

• 1897: JJ Thomson - Discovered negatively charged particles (electrons) through the Plum Pudding Model.

• 1911: Rutherford - Identified the nucleus via the Gold Foil Experiment.

• 1913: Bohr - Established energy levels for electrons resembling planetary orbits.

• 1920s: Various Scientists - Proposed electron cloud model.

• 1932: Chadwick - Discovered neutrons.

Identifying Atoms

• Protons

  • Positively charged particles located in the nucleus of an atom.

  • Mass: 1 atomic mass unit (a.m.u).

  • Determines the element through its atomic number (akin to a fingerprint for identification).

  • Rule: Atoms of the same element possess identical proton counts (atomic number).

• Isotopes

  • Variants of the same element differing by neutron count.

  • Example: 17O has 8 neutrons, while 18O has 10 neutrons.

• Mass Number

  • Represents the sum of protons and neutrons in an atom.

Interpreting the Periodic Table

Key Components:

1. Atomic Number

   • Located at the top: indicates the number of protons within the atom.

2. Atomic Symbol

   • One or two letters representing the element, derived from its name or Latin counterpart.

3. Atomic Mass

   • Average mass computed from isotopes of the element, given in atomic mass units (a.m.u).

Structure of the Atom

• Nucleus

  • Small center containing protons and neutrons.

  • Positively charged and extremely dense, contributing to the atom's mass.

Subatomic Particles

1. Protons

   • Charge: +

   • Mass: 1 a.m.u.

   • Location: Nucleus.

2. Neutrons

   • Charge: Neutral (0)

   • Mass: 1 a.m.u.

   • Location: Nucleus.

3. Electrons

   • Charge: -

   • Mass: approximately 0.00054 a.m.u.

   • Location: Electron cloud (around the nucleus).

Counting Subatomic Particles

• Counting Protons:

  • The number of protons defines the atomic number.

  • For neutral atoms, the number of electrons equals the number of protons.

• Mass Calculation

  • Atomic Mass

  • Mass Number = Number of Protons + Number of Neutrons (P + N = Mass Number).

Lesson 1 Check Questions

1. How does modern atomic theory differ from Dalton’s original theory of atoms?

2. What are the rules regarding subatomic particles and isotopes?

Lesson 2: The Periodic Table

Importance of Organization

• Elements must be structured to identify patterns and facilitate understanding.

Historical Development

• By 1869, there were 63 known elements that needed organizing.

• Dmitri Mendeleev developed a systematic arrangement based on chemical and physical properties as well as atomic mass.

Characteristics of the Periodic Table

• A grid reference containing all known elements.

• Arranged left to right and top to bottom by increasing atomic number, typically corresponding with atomic mass.

• As of 2016, features 118 confirmed elements (from Hydrogen [1] to Oganesson [118]).

Mendeleev’s Contributions

• Established a table predicting the properties of undiscovered elements.

• Insightful patterns within elements helped organize and predict chemical behavior.

Structure of the Table

1. Periods

   • Horizontal rows indicating the number of electron levels (shells).

   • Period number indicates the highest occupied energy level by valence electrons.

2. Groups

   • Vertical columns signifying the number of outermost “valence” electrons (1-8).

   • Elements within a group exhibit similar properties due to the equal number of valence electrons.

Divisions of the Periodic Table

Categories of Elements

1. Metals

   • Largest category located on the left side.

   • Characterized as solid, shiny, good electrical conductors.

   • Flexible (malleable) and can be drawn into wires (ductile).

   • Includes lanthanides and actinides.

2. Nonmetals

   • Fewer than metals, found on the right side.

   • Can be solid, liquid, or gas; poor conductors of heat and electricity.

3. Metalloids

   • Found between metals and nonmetals, possessing properties of both.

   • Commonly employed in semiconductor technology.

Groups of Elements

Metal Grouping

• Alkali Metals (Group 1)

  • Highly reactive and found in compounds; shiny and soft.

  • Reactivity increases down the group.

• Alkaline Earth Metals (Group 2)

  • Harder, denser, melting at higher temperatures than alkali metals.

  • Less reactive, never found in nature uncombined.

• Transition Metals (Groups 3-12)

  • Include common metals (e.g., Iron, Copper).

  • Usually hard and shiny, functioning as good conductors.

Nonmetals & Metalloids

• Boron Family (Group 13) - One metalloid (boron), with post-transition metals.

• Carbon Family (Group 14) - Carbon (nonmetal) accompanied by metalloids (silicon and germanium).

• Nitrogen Family (Group 15) - Contains a mix of nonmetals and metalloids.

• Oxygen Family (Group 16) - Composed of nonmetals and metalloids.

Unique Element Groups

• Halogens (Group 17)

  • Highly reactive nonmetals; salt formers, reactivity increases upwards in the group.

• Noble Gases (Group 18)

  • Inert gases that rarely form compounds due to full valence shells.

• Hydrogen

  • The most abundant element, lying at the top left corner; forms bonds rather than existing in a pure state.

Concept Check

• Label the periodic table considering the above information.