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Introduction to atoms and molecules. Chapter 2-1 – 2-5 Chapter 5-7 and 5-9 Chapter 4-5 – 4-6. Key concepts in this unit. Dalton’s atomic theory The electron, proton, and neutron Isotope, atomic number, mass number Molecules and molecular formulas Introduction to ions

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introduction to atoms and molecules

Introduction to atoms and molecules

Chapter 2-1 – 2-5

Chapter 5-7 and 5-9

Chapter 4-5 – 4-6

key concepts in this unit
Key concepts in this unit
  • Dalton’s atomic theory
  • The electron, proton, and neutron
  • Isotope, atomic number, mass number
  • Molecules and molecular formulas
  • Introduction to ions
  • Nomenclature for inorganic compounds
concept of the atom
Concept of the atom
  • Democritus (ancient Greek) --  (atomos); indivisible particle of a substance
  • Democritus’ ideas were promptly lost for the next ~ 2000 years or so.
dalton birth of the atomic theory 1803
Dalton—birth of the atomic theory (1803)

Dalton’s postulates:

  • Each element is composed of extremely small, indivisible particles called atoms.
  • All atoms of a given element are identical. Atoms of different elements are different, and have different properties, including mass.
dalton s atomic theory con t
Dalton’s atomic theory (con’t)
  • Atoms of an element are not changed into different types of atoms in a chemical reaction (i.e., you can’t turn lead into gold). Atoms are not created or destroyed in chemical reactions.
  • Compounds are formed when atoms of more than one element combine in small, whole number ratios.
  • A given compound always has the same relative number and kind of atoms.
dalton s theory explains laws
Dalton’s theory explains laws
  • law of constant composition
  • law of conservation of matter
  • law of multiple proportions
law of multiple proportions
Law of multiple proportions
  • Elements combining in more than one proportion will do so in multiples of whole numbers.
  • Examples:
    • H2O vs. H2O2.
    • NO NO2 N2O
fundamental particles
Fundamental particles
  • Masses of fundamental particles:
  • Atomic mass unit:

1 amu = 1.66053873  10-24 g

  • Atoms vary by atomic number (Z), which indicates ________________________.
  • Atoms contain the same number of electrons and protons (they have no charge).
  • Atomic number increases from left to right on periodic table.
  • The atomic weight of the atom is indicated below the symbol (in amu).
  • Why are these weights not whole numbers?
isotopes chap 5 7 5 9
Isotopes(chap 5-7, 5-9)
  • Atoms of an element have same number of protons in the nucleus.
  • Isotopes of an element vary in number of neutrons, but not in number of protons.
  • Atomic symbols: (“carbon-12”)
    • C: carbon
    • Subscript number: atomic number; number of protons
    • Superscript number: mass number, number of protons + neutrons
  • What is the atomic symbol for carbon-14?
average atomic weight
Average atomic weight
  • The average atomic weight (the one in the periodic table) is a weighted average.
  • Not all isotopes of an atom exist in the same amounts, and this must be taken into account.
Example 1: Let’s assume carbon has only 2 isotopes, 13C and 12C. The natural abundance of 13C is 1.0 %. What is the average atomic mass?
  • Example 2: Cl has 2 isotopes, 35C and 37C. The average atomic mass of Cl is 35.5 amu. What is the natural abundance of the two isotopes?
  • A molecule is ___________________________________________________________________.
    • In some cases, an atom is also a molecule. Example: He, Ne, Ar.
some elements exist in nature as molecules
Some elements exist in nature as molecules.
  • Diatomic molecular elements:
  • Polyatomic molecular elements:
  • Allotropes:
  • “shorthand” for writing and describing molecular composition.
    • Chemical formula:
    • Structural formula:
ionic compounds
Ionic compounds
  • Ionic compounds are composed of _________________________________.
    • Cation:
    • Anion:
  • Ionic compounds do not form molecular units, but are _______________ containing a defined ratio of cations and anions, forming a _________.
  • Formula unit:
The charge on an ion is designated by a number written above and to the right of the symbol
  • Examples:
    • Na1+ (or Na+)
    • Al3+
    • F-
    • O2-
periodic groups and ions
Periodic groups and ions
  • Normally (but not always) atoms in the following groups form ions of the designated charge.
  • This commonly occurs such that the # of electrons in the ion are equal to the # of electrons in the neighboring noble gas. We will discuss why this is so in quantum mechanics.
Transition metals often form more than one type of cation (i.e., different charges, such as Fe+, Fe2+, Fe3+.)
  • We have methods of writing formulas and names to indicate what charge the ions take.
identifying ionic compounds
Identifying ionic compounds
  • Generally (but again, not always) ionic compounds are formed by combining a metal and a nonmetal. Molecular compounds are generally formed by combining nonmetals.
  • The sum of the charges in an ionic compound must equal zero. The subscript of each ion indicates how many of that ion are in the formula.
  • Example: what is the formula for the ionic compound formed when Al3+ combines with O2-?
nomenclature rules
Nomenclature rules
  • Cations
  • Cations from metal atoms take the name of the metal.
    • Examples:
  • Cations from transition metals have the charge indicated by a Roman numeral.
    • Examples:
    • Note: there is an older method of naming transition metal cations. We will not use the older method in this course, but you should be aware of it in case you ever run into it. (p.141)
  • Cations from nonmetal atoms end in –ium.
    • Examples:

Table 2-3 (p. 55) and Table 4-11 (p. 142) contain the names and symbols of common cations that you should know.

  • monoatomic anions are named by replacing the ending of the element with –ide.
  • Examples:
polyatomic anions containing a variable number of oxygens. These occur in ternary salts.

The common anion is named using the suffix –ate.

The suffix or prefix of the name changes when the number of oxygens is increased or decreased

Note that the CHARGES on a “family” of oxyanions doesn’t change when the

number of oxygens varies. The oxidation number of the central atom does change.

We’ll discuss this later.

examples of ternary salts
Examples of ternary salts
  • sodium sulfate
  • iron (II) oxide
  • ammonium chloride
  • lithium carbonite
oxyanions with an H+ or 2 H+ added are indicated using the word hydrogen, or dihydrogen.
  • examples:

(Sometimes the addition of a hydrogen is indicated by the prefix bi-. For instance, HCO3- is the bicarbonate anion, as in NaHCO3, sodium bicarbonate. Use of either bi or hydrogen is correct).

Table 2-3 (p. 55) and Table 4-11 (p. 142) contain the names and symbols of common anions that you should know. You should also be able to properly name the relatives of the “common” oxyanions.

  • For the moment, an acid is an ionic compound that, when dissolved in water, yields a H+.

Double check: the formula for any ionic compound must = ZERO.

binary molecular compounds
Binary molecular compounds
  • The atom farthest to the left in the periodic table is usually written first.
  • For elements in the same group, the atom appearing lower in the table is usually written first.
  • the name of the second element in the formula name is given an –ide ending.
  • Greek prefixes indicate the number of each type of atom appearing in the molecule. (p. 142) You should know these prefixes.