Atoms, Isotopes, and More
ATOMS AND MOLECULES
SYMBOLS AND FORMULAS
Learning Objectives
Use symbols for chemical elements to write formulas for chemical compounds.
CHEMICAL SYMBOLS
Definition:
Atoms are the basic units of matter.
Elements are pure substances made of one type of atom.
Molecules are groups of two or more atoms bonded together.
Compounds are substances formed when two or more different elements chemically bond together.
Communication of identities:
There are 118 unique elements, each with its own unique chemical and physical properties.
Each element has a unique chemical symbol, usually based on its name, e.g.,
Calcium:
Symbol: Ca
Iron (from Latin):
Symbol: Fe
Reference:
Familiarity with the periodic table is essential for understanding chemical symbols.
Resource: Interactive periodic table from the Royal Society of Chemistry
URL: http://www.rsc.org/periodic-table
Available as a webpage or app.
COMPOUND FORMULAS
Definition:
Compound formulas are written using element symbols and the number of atoms of each element in a unit molecule.
Formula Structure:
Compound Formula = Symbol + Number of Atoms of that Element
Example:
Water:
Composed of two hydrogen atoms and one oxygen atom bonded together.
Formula: H2O
Interpretation:
2 Hydrogen (H) atoms and 1 Oxygen (O) atom in a water molecule.
LEARNING CHECK 2.1
Write Molecular Formulas for the Following Compounds:
Phosphoric acid:
Composition: 3 hydrogen (H) atoms, 1 phosphorus (P) atom, 4 oxygen (O) atoms
Formula: H3PO4
Sulfur trioxide:
Composition: 1 sulfur (S) atom, 3 oxygen (O) atoms
Formula: SO3
Glucose:
Composition: 6 carbon (C) atoms, 12 hydrogen (H) atoms, 6 oxygen (O) atoms
Formula: C6H12O6
INSIDE THE ATOM
Learning Objective
Identify the characteristics of protons, neutrons, and electrons.
WHAT’S IN AN ATOM?
Definition:
An atom is the smallest particle of matter produced by a chemical change, a limit of chemical subdivision for matter.
Subatomic Particles:
Fundamental subatomic particles include:
Proton
Neutron
Electron
FUNDAMENTAL SUBATOMIC PARTICLES
Size comparison:
Protons and neutrons are similar in size
Electrons are approximately 1800 times less massive than protons and neutrons.
Nucleus:
Protons and neutrons form the nucleus of an atom.
Characteristics required to know:
Symbols, charge, mass (u), location of each particle.
ISOTOPES
Learning Objective
Use the concepts of atomic number and mass number to determine the number of subatomic particles in isotopes and to write correct symbols for isotopes.
ATOMIC NUMBERS AND MASS NUMBERS
Definitions:
Atomic number (Z): the number of protons in an atom's nucleus.
All atoms of a particular element have the same number of protons (same atomic number).
Mass number (A): the sum of protons and neutrons in an atom's nucleus.
Isotopes:
Atoms with the same atomic number but different mass numbers due to differing numbers of neutrons.
Example: Hydrogen (H) has three isotopes, all having one proton but varying neutron counts.
Isotope Notation:
Written as: E (chemical symbol)
Alternatively: E-A (where A is the mass number).
LEARNING CHECK 2.2
Use the periodic table to answer the following questions about isotopes:
What are the atomic number, mass number, and isotope symbol for an atom that contains 4 protons and 5 neutrons?
Z = 4
A = 4 + 5 = 9
Be (Beryllium-9)
How many neutrons are contained in an atom of chlorine-37?
Calculation:
Number of Neutrons = A - Z = 37 - 17 = 20
How many protons and how many neutrons are contained in an atom with a mass number of 28 and the symbol Si?
Z = 14
Number of Neutrons = 28 - 14 = 14
Therefore, 14 protons and 14 neutrons.
RELATIVE MASSES OF ATOMS AND MOLECULES
Learning Objective
Use atomic weights of the elements to calculate molecular weights of compounds.
RELATIVE MASSES
Definition:
Atomic masses are minuscule and impractical for calculations. Relative masses represent relationships between the masses of different atoms using values listed under their atomic symbols on the periodic table.
Examples of atomic weights:
Nitrogen (N): 14.01 u
Carbon (C): 12.01 u
Oxygen (O): 16.00 u
Fluorine (F): 19.00 u
MASSES AND UNITS
Measurement:
Actual atomic masses can be measured by mass spectrometry and communicated in grams (g).
Atomic weight:
Mass of an average atom of an element expressed in atomic mass units (u).
Symbol: u, for atomic mass units.
Molecular weight: calculated by summing the atomic masses of the constituent atoms.
LEARNING CHECK 2.3
Relationships Between Atoms Using Atomic Weights:
How many nitrogen (N) atoms equal the mass of one iron (Fe) atom?
Nitrogen = 14.01 u, Iron = 55.85 u
Calculation:
4N = 1Fe => 55.85 / 14.01 = 3.986
How many carbon (C) atoms equal the mass of six helium (He) atoms?
Carbon = 12.01 u, Helium = 4.003 u
Calculation:
2C = 6He => 6(4.003) / 12.01 = 2.000
How many calcium (Ca) atoms equal the mass of two argon (Ar) atoms?
Calcium = 40.08 u, Argon = 39.95 u
Calculation:
2Ca = 2Ar => 2(39.95) / 40.08 = 1.994
LEARNING CHECK 2.4
Calculate Molecular Weights for the Following Compounds:
Sulfuric acid: H2SO4
MW = 2(at. wt. H) + 1(at. wt. S) + 4(at. wt. O)
MW = 2(1.008) + 1(32.06) + 4(16.00) = 98.076 = 98.08 u (4 significant figures)
Isopropyl alcohol (rubbing alcohol): C3H8O
MW = 3(at. wt. C) + 8(at. wt. H) + 1(at. wt. O)
MW = 3(12.01) + 8(1.008) + 1(16.00) = 60.094 = 60.09 u (4 significant figures)
Glucose: C6H12O6
MW = 6(at. wt. C) + 12(at. wt. H) + 6(at. wt. O)
MW = 6(12.01) + 12(1.008) + 6(16.00) = 180.156 = 180.2 u (4 significant figures)
ISOTOPES AND ATOMIC WEIGHTS
Learning Objective
Use isotope percent abundances and masses to calculate atomic weights of elements.
ATOMIC WEIGHTS OF ISOTOPES
Example of Sodium (Na):
Neutrons in a sodium atom:
Neutrons = atomic weight - protons = 23 - 11 = 12
Mass of the sodium nucleus:
Mass = (11 protons x 1) + (12 neutrons x 1) = 11 + 12 = 23 u
Symbol: 11 Na (atomic weight approx. 22.99)
Example of Gallium (Ga):
Gallium has 2 isotopes: Ga-69 (68.9257 u) and Ga-71 (70.9249 u).
Percent Abundance:
Gallium does not exist as a single isotope; instead, approx. 60% is Ga-69 and 40% is Ga-71.
Averaging Atomic Weights:
Calculate average atomic weight based on percent abundance.
Work through example 2.5 in the textbook and Learning Check 2.5 to practice percent abundance calculations!