Atoms

• Compound: A pure substance made up of two or more elements in a fixed ratio by mass
• Formula of a compound tells us the ratios of its constituent elements and identifies each element by its atomic symbol
• Examples of compound ratios:

  • NaCl: sodium atoms to chlorine atoms in sodium chloride is 1:1
  • H2O: hydrogen atoms to oxygen atoms in water is 2:1
• Element: A substance consisting of identical atoms

  • There are 118 known elements, with 98 occurring in nature
  • Element symbols consist of one or two letters
  • Names are derived from various sources like English names, important people in atomic science, geographic locations, planets, and mythological sources
• Mixture: A combination of two or more pure substances

  • Substances may be present in any mass ratio
  • Each substance in a mixture has different physical properties
  • Mixtures can be homogeneous or heterogeneous
  • Physical means can be used to separate mixtures into their component parts if the properties of the individual components are known
• Dalton's Atomic Theory by John Dalton (1766-1844):

  • All matter is composed of very tiny particles called atoms
  • Atoms of the same element have the same chemical properties
  • Compounds are formed by the chemical combination of two or more kinds of atoms
  • Molecules are tightly bound combinations of two or more atoms acting as a single unit
• Evidence for Dalton's Theory:

  • Law of Conservation of Mass: Matter can neither be created nor destroyed
  • Monatomic elements consist of single atoms (e.g., helium, neon)
  • Diatomic elements occur as molecules of two atoms (e.g., H2, N2)
  • Polyatomic elements have three or more atoms per molecule (e.g., O3, P4, S8)
• Subatomic Particles:

  • Mass is given in atomic mass units (amu)
  • 1 amu is the mass of an atom of carbon with 6 protons and 6 neutrons
  • Relative sizes of the atomic nucleus and an atom can vary
• Mass and Atomic Number:

  • Mass number is the sum of protons and neutrons in an atom's nucleus
  • Atomic number is the number of protons in an atom's nucleus
  • Isotopes are atoms with the same number of protons but different numbers of neutrons
• Atomic Weight:

  • Atomic weight is the weighted average of the masses of naturally occurring isotopes of an element
• Mass and Size of an Atom:

  • Atoms have a concentrated mass in the nucleus
  • The diameter and density of an atom's nucleus can vary
• Periodic Table by Dmitri Mendeleyev (1834-1907):

  • Elements are arranged in order of increasing atomic weight
  • Certain properties recur periodically when elements are arranged this way
  • Elements with recurring properties are placed in the same column
• Metals:

  • Solids at room temperature (except for Hg)
  • Shiny, conduct electricity, ductile, and malleable
  • Form alloys (e.g., brass is an alloy of copper and zinc)
  • Tend to give up electrons in chemical reactions
• Nonmetals:

  • Except for hydrogen (H), located on the right side of the Periodic Table
  • Except for graphite, do not conduct electricity
  • Tend to accept electrons in chemical reactions
• Metalloids:

  • Have properties of both metals and nonmetals
  • Shiny like metals but do not conduct electricity
  • Examples include boron, silicon, germanium, arsenic, antimony, and tellurium
  • Silicon is a semiconductor, not conducting electricity under certain voltages but becoming a conductor at higher voltages
• Electron Configuration:

  • Arrangement of electrons in the extranuclear space
  • Electrons in an atom have quantized energy levels
  • Ground-state electron configuration is the lowest energy state of an atom
• Rules for Electron Configuration:

  1. Orbitals fill in order of increasing energy from lowest to highest

  2. Each orbital can hold up to two electrons with paired spins

  3. Orbitals of equal energy are half-filled before completely filled
• Orbital Box Diagrams:

  • A box represents an orbital
  • An arrow represents an electron
  • Pair of arrows with opposite directions represents paired electrons with paired spins
• Noble Gas Notation:

  • Indicates electron configuration of all filled shells
  • Uses the symbol of the noble gas preceding the atom
• Valence Shell and Valence Electron:

  • Valence shell is the outermost incomplete shell
  • Valence electron is an electron in the valence shell
• Periodic Property:

  • Periodic Table constructed based on trends in chemical properties
  • Periodicity in electron configuration explains periodicity in chemical properties
  • Elements in the same group have the same configuration in their outer shells
• Atomic Size:

  • Determined by the radius of outermost occupied orbitals
  • Example: Chlorine atom size determined by three 3p orbitals
• Ionization Energy:

  • Energy required to remove the most loosely held electron from an atom in the gaseous state
  • Lithium losing an electron becomes a lithium ion with a positive charge
• The nucleus is the central part of an atom, containing protons and neutrons.
• Protons are positively charged particles found in the nucleus.
• Electrons are negatively charged particles that surround the nucleus at different energy levels called orbits or shells.
• Isotopes are atoms of the same element but with different numbers of neutrons.
• Neutrons are neutral particles also located within the nucleus.
• Atomic number (Z) refers to the number of protons present in an atom's nucleus.
• Relative atomic mass (Mr) indicates the average relative mass of all naturally occurring isotopes of an element.
• Mass number (A) represents the total number of protons and neutrons in an atom's nucleus.
• Nuclear reactions involve changes in the number of protons and/or neutrons within the nucleus.
• The periodic table arranges elements based on their atomic structure, including atomic number and electron configuration.
• Elements can be classified as metals, non-metals, metalloids, noble gases, halogens, alkali metals, alkaline earth metals, transition metals, lanthanides, actinides, and inner transition metals.
• Metallic bonding occurs when metal atoms lose electrons from their outermost orbitals and form positive ions held together by electrostatic attraction.
• Metallic bonding occurs when positively charged metal ions attract delocalized electrons, resulting in strong intermolecular forces between neighboring positive ions.
• Electron configurations describe how electrons are arranged around the nucleus of an atom.
• Non-metal compounds have covalent bonds formed through sharing electrons.
• Non-metal bonds occur through covalent or ionic bonding, with shared pairs of electrons forming molecules or compounds.
• Ionic compounds are formed through ionic bonding, where one or more electrons are transferred from one element to another during chemical reaction.