Atoms and Elements Study Guide

Elements and Symbols

• Definition of Elements

  • Elements are pure substances from which all other things are built.
  • They consist of matter that cannot be broken down into simpler substances.
  • Elements are listed on the inside front cover of the chemistry text.
  • Examples of elements include gold, carbon, and aluminum.

• Origins of Element Names

  • Element names are derived from various sources, including:
    • Planets.
    • Mythological figures.
    • Minerals.
    • Colors.
    • Geographic locations.
    • Famous people.

• Chemical Symbols

  • A chemical symbol represents the name of an element.
  • Symbols consist of one to two letters.
  • The first letter is always capitalized.
  • One-Letter Symbols:
    • Carbon: $C$
    • Nitrogen: $N$
    • Fluorine: $F$
    • Oxygen: $O$
  • Two-Letter Symbols:
    • Cobalt: $Co$
    • Calcium: $Ca$
    • Aluminum: $Al$
    • Magnesium: $Mg$

• Symbols Derived from Latin Names

  • Several symbols come from the element's original Latin name:
    • Copper: $Cu$ (from cuprum)
    • Gold: $Au$ (from aurum)
    • Iron: $Fe$ (from ferrum)
    • Silver: $Ag$ (from argentum)

• Chemistry Link to Health: Toxicity of Mercury

  • Mercury ($Hg$) is a silvery, shiny element that remains a liquid at room temperature.
  • Methods of Entry: Mercury enters the body through mercury vapor inhalation, skin contact, or the ingestion of contaminated water or food.
  • Toxicity: It is toxic to both the central and peripheral nervous systems.
  • Biological Impact: Once inside the body, mercury destroys proteins and disrupts cell function.
  • Effects of Long-term Exposure:
    • Damage to the brain and kidneys.
    • Mental retardation.
    • Decreased physical development.
  • Sources of Contamination:
    • Industrial wastes.
    • Fish and seafood.
    • Batteries.
    • Compact fluorescent bulbs.

The Periodic Table

• Organization of the Periodic Table

  • The table organizes $118$ elements.
  • Elements are placed in order of increasing atomic mass and arranged into groups based on similar properties.
  • Horizontal Rows (Periods): Horizontal rows of elements are counted from top to bottom as Periods $1-7$.
  • Vertical Columns (Groups): Groups contain elements with similar properties in vertical columns.

• Group Numbers and Classification

  • Group numbers are written at the top of each vertical column.
  • Letter System:
    • Letter $A$ is used for representative elements (Groups $1A-8A$).
    • Letter $B$ is used for transition elements (Groups $3B-12B$).
  • Numbered System: An alternative system uses numbers $1-18$ across the table from left to right.

• Specific Group Names

  • Group $1A(1)$: Alkali metals (includes Lithium ($Li$), Sodium ($Na$), Potassium ($K$), Rubidium ($Rb$), and Cesium ($Cs$)).
  • Group $2A(2)$: Alkaline earth metals. These are shiny but less reactive than Group $1A$; includes Beryllium ($Be$), Magnesium ($Mg$), Calcium ($Ca$), Strontium ($Sr$), Barium ($Ba$), and Radium ($Ra$).
  • Group $7A(17)$: Halogens (includes Fluorine ($F$), Chlorine ($Cl$), Bromine ($Br$), and Iodine ($I$)).
  • Group $8A(18)$: Noble gases.

• Classification by Properties: Metals, Nonmetals, and Metalloids

  • A heavy zigzag line separates the metals and nonmetals on the table.
  • Metals:
    • Located to the left of the zigzag line (except for hydrogen).
    • Characteristics: Shiny, ductile, good conductors of heat and electricity.
    • Physical state: Solids, except for mercury ($Hg$), which is a liquid.
  • Nonmetals:
    • Located to the right of the zigzag line.
    • Characteristics: Dull, brittle, poor conductors of heat and electricity, but good insulators.
    • Physical properties: Low densities and low melting points.
  • Metalloids:
    • Located along the heavy zigzag line.
    • Characteristics: Exhibit properties of both metals and nonmetals.
    • Conductivity: Better than nonmetals but not as good as metals.
    • Usage: Used as semiconductors and insulators (can be modified to function as either).

• Chemistry Link to Health: Elements Essential to Health

  • Vital Elements: $20$ elements are essential for the survival and well-being of the human body.
  • Primary Elements: Four elements—oxygen ($O$), carbon ($C$), hydrogen ($H$), and nitrogen ($N$—make up $96\%$ of total body mass.
  • Water Content: Most hydrogen and oxygen is found in water, which accounts for $55\%$ to $60\%$ of body mass.
  • Macrominerals: $Ca, P, K, Cl, S, Na, Mg$.
    • Involved in bone and teeth formation.
    • Maintenance of the heart, blood vessels, muscle contraction, and nerve impulses.
    • Regulation of acid-base balance and cellular metabolism.

The Atom

• Definition of the Atom

  • An atom is the smallest particle of an element that retains the characteristics of that element.

• Dalton’s Atomic Theory

  • Atoms are tiny particles of matter.
  • Atoms of a specific element are similar to each other and different from atoms of other elements.
  • Atoms of two or more different elements combine to form compounds.
  • Chemical reactions involve the rearrangement of atoms to form new combinations.
  • Atoms are never created or destroyed during a chemical reaction.

• Subatomic Particles and Electrical Charges

  • Protons: Have a positive ($+$) charge.
  • Electrons: Have a negative ($-–$) charge.
  • Neutrons: Have no charge (neutral).

• Models of the Atom

  • Thomson’s "Plum-Pudding" Model: Proposed that protons and electrons were randomly distributed in a positively charged cloud, similar to plums in a pudding.
  • Rutherford’s Gold Foil Experiment: Concluded that there must be a small, dense, positively charged nucleus in the center of the atom that deflects positive particles.

• Structure of the Atom

  • Nucleus: Located in the center; contains protons and neutrons. It represents most of the atom’s mass.
  • Electron Space: Electrons occupy a large, mostly empty space surrounding the nucleus. While electrons account for the atom’s volume, the protons and neutrons contribute almost all the mass in a tiny volume.

• Mass of the Atom

  • Atomic Mass Unit (amu): Used to measure the mass of subatomic particles.
  • Definition: $1\,amu$ is equal to $1/12$ of the mass of a Carbon-$12$ atom (which has $6$ protons and $6$ neutrons).
  • Biology connection: $1\,amu$ is equivalent to $1\,Dalton\,(Da)$.
  • Mass of Electrons: Electrons have such a small mass that they are excluded from the total mass calculation of an atom.

Atomic Number and Mass Number

• The Atomic Number

  • The atomic number is a whole number specific to each element.
  • It is the same for all atoms of a specific element.
  • It is equal to the number of protons in an atom.
  • It appears above the chemical symbol of an element (e.g., Sodium ($Na$) has an atomic number of $11$).

• Neutral Atoms

  • For neutral atoms, the net charge is zero.
  • This requires that the number of protons ($+$) equals the number of electrons ($-–$).
  • Example: Aluminum ($Al$) has $13$ protons and $13$ electrons; $13+ + 13- = 0$.

• The Mass Number

  • Represents the total number of particles (nucleons) in the nucleus.
  • $\text{Mass Number} = \text{number of protons} + \text{number of neutrons}$.
  • It is always a whole number and does not appear on the periodic table.
  • To find neutrons: $\text{Number of Neutrons} = \text{Mass Number} - \text{Atomic Number}$.

Isotopes and Atomic Mass

• Isotopes

  • Isotopes are atoms of the same element that have different mass numbers.
  • They have the same number of protons but different numbers of neutrons.
  • Atomic symbols for isotopes (e.g., ${}_{12}^{24}Mg$) distinguish between them.

• Example: Magnesium ($Mg$) Isotopes

  • $Mg-24$: $12$ protons, $12$ neutrons, $12$ electrons, mass $23.99\,amu$, abundance $78.70\%$.
  • $Mg-25$: $12$ protons, $13$ neutrons, $12$ electrons, mass $24.99\,amu$, abundance $10.13\%$.
  • $Mg-26$: $12$ protons, $14$ neutrons, $12$ electrons, mass $25.98\,amu$, abundance $11.17\%$.

• Calculating Atomic Mass

  • Atomic mass is the weighted average of all naturally occurring isotopes of an element.
  • Calculation Steps:
    1. Divide each isotope's percent abundance by $100$ to get the fractional abundance.
    2. Multiply the mass of each isotope by its fractional abundance.
    3. Sum the contributions of each isotope.
  • Example: For Magnesium, the weighted average is $24.31\,amu$.

Electron Energy Levels

• Electromagnetic Radiation

  • Includes light, rainbow colors, and X-rays.
  • Consists of energy particles moving as waves.
  • Wavelength: The distance between wave peaks.
  • High-energy radiation has shorter wavelengths; low-energy radiation has longer wavelengths.

• Atomic Spectrum

  • When light from a heated element passes through a prism, it separates into distinct lines of color and dark areas.
  • Each element has a unique atomic spectrum.

• Energy Levels

  • Energy levels are assigned principal quantum numbers ($n$) starting at $n = 1, n = 2, \dots$.
  • Energy increases as $n$ increases (as electrons get farther from the nucleus).
  • Energy is quantized, meaning electrons can only exist at specific energy values.
  • Transitions:
    • Electrons move to higher levels by absorbing energy.
    • Electrons emit light (a photon) when they fall back to a lower energy level.
    • Energy change ($E$) equals the difference between the two energy levels.

• Sublevels and Orbitals

  • Each energy level contains one or more sublevels.
  • Number of sublevels in a level = the value of $n$.
  • Sublevels are designated as $s, p, d, f$.
  • Order of energy: $s < p < d < f$.
  • Orbital Types:
    • $s$ Orbital: Spherical shape. The size increases with higher energy levels. Each $s$ sublevel has $1$ orbital ($2$ electrons).
    • $p$ Orbital: "Dumbbell" shape (two lobes), starting at $n = 2$. Each $p$ sublevel has $3$ orbitals ($6$ electrons) arranged along $x, y,$ and $z$ axes.
    • $d$ Orbital: Most consist of four lobes; one consists of two lobes and a ring. Each $d$ sublevel has $5$ orbitals ($10$ electrons).
    • $f$ Orbital: Each $f$ sublevel has $7$ orbitals ($14$ electrons).
  • Pauli Exclusion Principle: Each orbital holds a maximum of two electrons with opposite spins.

Electron Configurations

• Notation

  • Orbital Diagrams: Use boxes to represent orbitals and arrows for electrons.
  • Rules for Filling:
    1. Fill orbitals from lowest to highest energy level.
    2. Fill orbitals within the same sublevel one at a time before pairing electrons.
  • Electron Configuration: Indicates the placement of electrons (e.g., Carbon: $1s^2 2s^2 2p^2$).
  • Abbreviated Form: Uses the previous Noble Gas in brackets (e.g., Lithium ($Li$): $[He] 2s^1$).

• Periodic Table Blocks

  • $s$ block: Groups $1A$ and $2A$.
  • $p$ block: Groups $3A$ to $8A$.
  • $d$ block: Transition elements (starts after Calcium, atomic number $20$).
  • $f$ block: Inner transition elements (the two rows at the bottom).

Trends in Periodic Properties

• Valence Electrons

  • These are the electrons in the outermost energy level.
  • For representative elements, the group number identifies the number of valence electrons.
  • Lewis Symbols (Electron-dot symbols): Represent valence electrons as dots around the element's symbol. One to four electrons are single dots; five to eight involve pairs.

• Atomic Size

  • Determined by atomic radius (distance between nucleus and outermost electrons).
  • Trend: Increases from top to bottom of a group; decreases from left to right across a period (due to increased proton pull).

• Ionization Energy

  • The energy required to remove one of the outermost electrons.
  • Equation: $Na(g) + \text{energy} \rightarrow Na^+(g) + e^-$.
  • Trend: Decreases down a group; increases across a period from left to right.

• Metallic Character

  • Refers to the ease with which an element loses valence electrons.
  • Trend: Increases going down a group; decreases from left to right across a period.