Chapter 4 Notes: Atoms and Elements

Atoms and Elements

An atom is the smallest identifiable unit of an element.
An element is a substance that cannot be broken down into simpler substances.
There are about 91 different elements in nature.
Scientists have succeeded in making about 20 synthetic elements (not found in nature).

The Atomic Theory

Democritus suggested that if you divide matter into smaller pieces, you end up with tiny, indestructible particles.
Democritus called them atomos, or “atoms,” meaning “indivisible.”
Democritus is the first person on record to have postulated that matter was composed of atoms.

Dalton's Atomic Theory (1808)

John Dalton formalized a theory of atoms that gained broad acceptance. Dalton’s atomic theory has three parts:

Each element is composed of tiny indestructible particles called atoms.
All atoms of a given element have the same properties that distinguish them from the atoms of other elements.
Atoms combine in simple, whole-number ratios to form compounds.

Modern Evidence for the Atomic Theory

Scientists at IBM used a special microscope, called a scanning tunneling microscope (STM), to move xenon atoms to form the letters I, B, and M.
The cone shape of these atoms is due to the peculiarities of the instrumentation. Atoms are, in general, spherical in shape.

Discovery of the Electron

J. J. Thomson (1856–1940) discovered a smaller and more fundamental particle called the electron.
Electrons are negatively charged.
Electrons are much smaller and lighter than atoms.
Electrons are uniformly present in many different kinds of substances.
Atoms must contain positive charge that balances the negative charge of electrons.

Plum Pudding Model

In the model suggested by J. J. Thomson, negatively charged electrons were held in a sphere of positive charge.

Rutherford's Gold Foil Experiment

Tiny particles called alpha-particles were directed at a thin sheet of gold foil.
Most particles passed right through the foil.
A few particles were deflected at large angles.

Discovery of the Atomic Nucleus

Expected Result (Plum Pudding Model): Alpha-particles would pass right through the gold foil with minimal deflection.
Actual Result: A small number of alpha-particles were deflected or bounced back.

Rutherford's Nuclear Theory

Most of the atom’s mass and all of its positive charge are contained in a small core called the nucleus.
Most of the volume of the atom is empty space through which the tiny, negatively charged electrons are dispersed.
The number of negatively charged electrons outside the nucleus is equal to the number of positively charged particles (protons) inside the nucleus, so that the atom is electrically neutral.

Nuclear Model

Volume of atom is mostly empty space.

Properties of Protons, Neutrons, and Electrons

Protons and neutrons have very similar masses.
Electrons have almost negligible mass.

Atomic Number (Z)

The number of protons in the nucleus of an atom identifies it as a particular element.
The number of protons in the nucleus of an atom is its atomic number and is given the symbol Z.

Periodic Table of Elements

The periodic table of the elements lists all known elements according to their atomic numbers.

Origins of the Names of the Elements

Most chemical symbols are based on the English name of the element.
Some symbols are based on Latin names:
- Lead: $Pb$ (plumbum)
- Mercury: $Hg$ (hydrargyrum)
- Iron: $Fe$ (ferrum)
- Silver: $Ag$ (argentum)
- Tin: $Sn$ (stannum)
- Copper: $Cu$ (cuprum)
- Potassium: $K$ (kalium)
- Sodium: $Na$ (natrium)
Curium is named after Marie Curie.
Bromine originates from the Greek word bromos, meaning “stench.”

The Periodic Law and the Periodic Table

Mendeleev proposed that when the elements are arranged in order of increasing relative mass, certain sets of properties recur periodically.
Arrange elements in rows so that similar properties align in the same vertical columns.

Classification of Elements

The elements in the periodic table can be broadly classified as metals, nonmetals, and metalloids.

Metals

Metals occupy the left side of the periodic table.
Metals are good conductors of heat and electricity.
Metals can be pounded into flat sheets (malleability).
Metals can be drawn into wires (ductility).
Metals are often shiny (lustrous).
Metals tend to lose electrons when they undergo chemical changes.
Examples: iron, magnesium, chromium, and sodium.

Nonmetals

Nonmetals occupy the upper right side of the periodic table.
Nonmetals have more varied properties; some are solids at room temperature, while others are gases.
They tend to be poor conductors of heat and electricity.
They gain electrons when they undergo chemical changes.
Examples: oxygen, nitrogen, chlorine, and iodine.

Metalloids

Metalloids lie along the zigzag diagonal line dividing metals and nonmetals.
Metalloids, also called semimetals, display mixed properties.
Metalloids are also called semiconductors because of their intermediate electrical conductivity.
Semiconductors useful in the manufacture of electronic devices that are central to computers and cell phones.
Examples: Silicon, arsenic, and germanium.

Groups of Elements

Alkali Metals

The alkali metals include lithium, sodium, potassium, rubidium, and cesium.

Alkaline Earth Metals

The alkaline earth metals include beryllium, magnesium, calcium, strontium, and barium.

Halogens

The halogens include fluorine, chlorine, bromine, iodine, and astatine.

Noble Gases

The noble gases include helium, neon, argon, krypton, and xenon.

Diatomic Molecules

Diatomic molecules each contain exactly two atoms.
There are 7 diatomic elements.

Ions: Losing and Gaining Electrons

In chemical reactions, atoms often lose or gain electrons to form charged particles called ions.
Positive ions are called cations.
Negative ions are called anions.
Ion charges are usually written with the magnitude of the charge first followed by the sign of the charge (e.g., $3+$ , not $+3$ ).

Forming Cations (Losing Electrons)

In reactions, lithium atoms lose one electron ( $e^−$ ) to form $Li^+$ ions.
The charge of an ion depends on how many electrons were gained or lost.
For the $Li^+$ ion with 3 protons and 2 electrons, the charge is $3 − 2 = 1+$ .

Forming Anions (Gaining Electrons)

In reactions, fluorine atoms gain 1 electron to form $F^−$ ions.
The charge of an ion depends on how many electrons were gained or lost and is given by the formula.
For the $F^−$ ion with 9 protons and 10 electrons, the charge is $9 − 10 = 1-$ .

Predicting Ion Charges

The key to predicting the charge acquired by an element is its position in the periodic table relative to the noble gases.
The number associated with the letter A above each main-group column in the periodic table gives the number of valence electrons for the elements in that column.
Main-group elements tend to form ions that have the same number of valence electrons as the nearest noble gas.

Isotopes: When the Number of Neutrons Varies

All atoms of a given element have the same number of protons.
They do not necessarily have the same number of neutrons.
Atoms with the same number of protons but different numbers of neutrons are called isotopes.

Isotope Symbols

The mass number (A) is the sum of the number of protons and the number of neutrons.
The number of neutrons in an isotope is the difference between the mass number and the atomic number.

Natural Abundance of Isotopes

Naturally occurring neon contains three different isotopes: Ne-20, Ne-21, and Ne-22.

Atomic Mass: The Average Mass of an Element’s Atoms

The atomic mass of each element listed in the periodic table represents the average mass of the atoms that compose that element.
Naturally occurring chlorine consists of 75.77% chlorine-35 (mass 34.97 amu) and 24.23% chlorine-37 (mass 36.97 amu).

Calculating Atomic Mass

In general, atomic mass is calculated according to the following equation:

$Atomic \ mass = (Fraction \ of \ isotope \ 1 × Mass \ of \ isotope \ 1) + (Fraction \ of \ isotope \ 2 × Mass \ of \ isotope \ 2) + (Fraction \ of \ isotope \ 3 × Mass \ of \ isotope \ 3) + …$

where the fractions of each isotope are the percent natural abundances converted to their decimal values.

Example: Calculating Atomic Mass of Gallium

Gallium has two naturally occurring isotopes: Ga-69 (mass 68.9256 amu, abundance 60.11%) and Ga-71 (mass 70.9247 amu, abundance 39.89%).
Convert the percent natural abundances into decimal form by dividing by 100.