Chemistry: Atoms and Elements

Atoms and Elements

Graphite and Atoms

• If graphite is cut into smaller and smaller pieces, it eventually yields individual carbon atoms.
• The word "atom" comes from the Greek "atomos," meaning "indivisible."
• Carbon atoms cannot be divided into smaller pieces while still retaining their carbon identity.
• Atoms compose all ordinary matter.

Brownian Motion and Confirmation of Atoms

• Robert Brown (1773–1858) observed continuous motion of water-suspended particles from pollen grains under a microscope.
• Albert Einstein (1879–1955) developed a theory in 1905 that quantitatively explained Brownian motion.
• Jean Perrin (1870–1942) confirmed Einstein’s model through measurements in 1908 and was awarded the Nobel Prize in Physics in 1926.
• The work of Einstein and Perrin eliminated lingering doubts about the particulate nature of matter.

Imaging and Moving Individual Atoms

• Atoms are pivotal in connecting macroscopic and microscopic worlds.
• An atom is the smallest identifiable unit of an element.
• There are about 91 naturally occurring elements and over 20 synthetic elements.

Early Ideas About Matter

• Leucippus (5th century B.C.) and Democritus (460–370 B.C.) proposed that matter is composed of small, indestructible particles.
• Democritus stated, "Nothing exists except atoms and empty space; everything else is opinion."
• They suggested that atoms existed in different shapes and sizes, moving randomly through empty space.
• Plato and Aristotle opposed atomic ideas, believing matter had no smallest parts and that substances were composed of fire, air, earth, and water.

Modern Atomic Theory and Laws

• John Dalton (1766–1844) provided evidence supporting early atomic ideas.
• The atomic theory grew out of observations and laws.
• Three important laws leading to the development of atomic theory include:
  • Law of conservation of mass
  • Law of definite proportions
  • Law of multiple proportions

Law of Conservation of Mass

• Antoine Lavoisier formulated the law of conservation of mass.
• The law states that in a chemical reaction, matter is neither created nor destroyed; the total mass of substances remains unchanged.
• This law aligns with the concept that matter consists of small, indestructible particles.

Law of Definite Proportions

• Joseph Proust made observations on compound composition in 1797.
• The law of definite proportions states that all samples of a compound have the same proportions of constituent elements, regardless of source or preparation.
• Also known as the law of constant composition.
• For example, the decomposition of 18.0 g of water yields 16.0 g of oxygen and 2.0 g of hydrogen, with an oxygen-to-hydrogen mass ratio of \frac{16.0 \text{ g O}}{2.0 \text{ g H}} = 8.0 , or 8:1.

Law of Multiple Proportions

• John Dalton published the law of multiple proportions in 1804.
• When two elements (A and B) form two different compounds, the masses of element B that combine with 1 g of element A can be expressed as a ratio of small whole numbers.
• An atom of A combines with one, two, or three atoms of B (e.g., AB, AB2, AB3).
• Carbon monoxide and carbon dioxide contain the same elements (carbon and oxygen).
  • In carbon dioxide, the mass ratio of oxygen to carbon is 2.67:1 (2.67 g O per 1 g C).
  • In carbon monoxide, the mass ratio of oxygen to carbon is 1.33:1 (1.33 g O per 1 g C).
• The ratio of oxygen mass per 1 g carbon in carbon dioxide to that in carbon monoxide is a small whole number: \frac{2.67}{1.33} = 2 .

Dalton’s Atomic Theory

• Dalton’s atomic theory explained the laws as follows:
  1. Each element is composed of tiny, indestructible particles called atoms.
  2. All atoms of a given element have the same mass and properties, distinguishing them from atoms of other elements.
  3. Atoms combine in simple, whole-number ratios to form compounds.
  4. Atoms of one element cannot change into atoms of another element; in chemical reactions, atoms only change how they are bound together.

Discovery of the Electron

• J. J. Thomson (1856–1940) conducted cathode ray experiments using a partially evacuated glass tube (cathode ray tube).
• He discovered that a beam of particles, called cathode rays, traveled from the negatively charged electrode (cathode) to the positively charged one (anode).
• Properties of cathode ray particles:
  • Travel in straight lines
  • Independent of the composition of the material from which they originate
  • Carry a negative electrical charge
• Thomson measured the charge-to-mass ratio of cathode ray particles using electric and magnetic fields.
• The measured value was 1.76 \times 10^8 \frac{\text{C}}{\text{g}}.
• J. J. Thomson discovered the electron, a negatively charged, low-mass particle present within all atoms.

Millikan’s Oil Drop Experiment

• Robert Millikan (1868–1953) deduced the charge of a single electron.
• By measuring the electric field strength required to halt the free fall of oil drops, Millikan calculated the charge of each drop.
• The measured charge on any drop was always a whole-number multiple of 1.60 \times 10^{-19} \text{ C}, the fundamental charge of a single electron.
• Knowing Thomson’s mass-to-charge ratio, the mass of an electron was deduced:
  • \text{mass} = \frac{\text{charge}}{\text{charge/mass}}

Structure of the Atom

• J. J. Thomson proposed that negatively charged electrons were small particles held within a positively charged sphere (plum-pudding model).
• Ernest Rutherford (1871–1937) conducted the gold foil experiment in 1909, attempting to confirm Thomson’s model.

Rutherford’s Gold Foil Experiment

• Rutherford directed positively charged alpha particles at an ultra-thin sheet of gold foil.
• Most particles passed through the foil, but some were deflected, and approximately 1 in 20,000 bounced back.
• Rutherford concluded that matter contains large regions of empty space dotted with small regions of very dense matter.
• He proposed the nuclear theory of the atom:
  1. Most of the atom’s mass and positive charge are in a small core called the nucleus.
  2. Most of the atom’s volume is empty space, throughout which tiny, negatively charged electrons are dispersed.
  3. There are as many negatively charged electrons outside the nucleus as positively charged particles (protons) within the nucleus, making the atom electrically neutral.

Neutrons

• Rutherford’s model was incomplete.
• James Chadwick (1891–1974) demonstrated the existence of neutrons, neutral particles within the nucleus.
• The mass of a neutron is similar to that of a proton, but it has no electrical charge.
• The helium atom is four times as massive as the hydrogen atom because it contains two protons and two neutrons, while hydrogen contains only one proton and no neutrons.

Subatomic Particles

• All atoms are composed of:
  • Protons
  • Neutrons
  • Electrons
• Protons and neutrons have nearly identical masses:
  • Mass of proton: 1.67262 \times 10^{-27} \text{ kg}
  • Mass of neutron: 1.67493 \times 10^{-27} \text{ kg}
  • Mass of electron: 0.00091\times 10^{-27} \text{ kg}
• The charge of the proton and electron are equal in magnitude but opposite in sign.
• Neutrons have no charge.

Elements and Atomic Number

• The number of protons in its nucleus defines the identity of an atom and is its atomic number (Z).
• Each element is identified by a unique atomic number and chemical symbol.
• Chemical symbols are one- or two-letter abbreviations on the periodic table.
  • Helium: He
  • Carbon: C
  • Nitrogen: N

Isotopes

• Atoms of a given element have the same number of protons but may have different numbers of neutrons.
• Atoms with the same number of protons but different numbers of neutrons are called isotopes.
• For example, neon atoms contain 10 protons but may contain 10, 11, or 12 neutrons.
• The relative amount of each isotope in a naturally occurring sample is roughly constant.
• These percentages are called the natural abundance of the isotopes.
• Advances in mass spectrometry have allowed accurate measurements revealing small variations in natural abundance.
• The sum of the number of neutrons and protons in an atom is its mass number (A).
  • A = \text{number of protons (p)} + \text{number of neutrons (n)}
• Notation for isotopes:
  • ^{20}{10}Ne, ^{21}{10}Ne, ^{22}_{10}Ne

Ions

• The number of electrons in a neutral atom is equal to the number of protons (atomic number Z).
• Atoms can lose or gain electrons during chemical changes, becoming charged particles called ions.
  • Positively charged ions (e.g., Na^+) are called cations.
  • Negatively charged ions (e.g., F^−) are called anions.

The Periodic Law and The Periodic Table

• In 1869, Mendeleev noticed that certain groups of elements had similar properties and when elements are listed in order of increasing mass, these similar properties recurred periodically.
• Periodic means exhibiting a repeating pattern.
• The periodic law states that when elements are arranged in order of increasing mass, certain sets of properties recur periodically.
• Mendeleev organized known elements in a table, arranging rows so that elements with similar properties fall in the same vertical columns.
• Mendeleev’s table contained gaps, allowing him to predict the existence and properties of undiscovered elements (e.g., eka-silicon, later discovered as germanium).
• In the modern table, elements are listed in order of increasing atomic number rather than relative mass.

Classification of Elements

• Elements are classified as:
  • Metals
  • Nonmetals
  • Metalloids

Metals

• Located on the lower-left side and middle of the periodic table.
• Properties:
  • Good conductors of heat and electricity
  • Malleable (can be pounded into flat sheets)
  • Ductile (can be drawn into wires)
  • Often shiny
  • Tend to lose electrons during chemical changes
• Examples: Chromium, copper, strontium, lead.

Nonmetals

• Located on the upper-right side of the periodic table.
• Varied properties:
  • Some are solids (C, P, S, Se, I)
  • One is a liquid (Br)
  • Eleven are gases (H, He, N, O, F, Ne, Cl, Ar, Kr, Xe, Rn)
• Properties:
  • Poor conductors of heat and electricity
  • Not ductile or malleable
  • Gain electrons during chemical changes
• Examples: Oxygen, carbon, sulfur, bromine, iodine.

Metalloids

• Also called semimetals.
• Lie along the zigzag diagonal line dividing metals and nonmetals.
• Exhibit mixed properties.
• Several are semiconductors with intermediate, temperature-dependent electrical conductivity.

Main-Group and Transition Elements

• Main-group elements have predictable properties based on their periodic table position.
• Transition elements (or transition metals) have less predictable properties based simply on their position.

Groups and Periods

• The periodic table has vertical columns (groups or families) and horizontal rows (periods).
• There are 18 groups and 7 periods.
• Groups are numbered 1–18 (or using A and B designations).
• Main-group elements are in columns labeled with the letter A (1A–8A or groups 1, 2, and 13–18).
• Transition elements are in columns labeled with the letter B (or groups 3–12).

Special Groups

• Noble Gases (Group 8A): Mostly unreactive; e.g., helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe).
• Alkali Metals (Group 1A): Reactive metals; e.g., lithium (Li), sodium (Na), potassium (K), rubidium (Rb).
• Alkaline Earth Metals (Group 2A): Fairly reactive metals; e.g., calcium (Ca), magnesium (Mg), strontium (Sr), barium (Ba).
• Halogens (Group 7A): Very reactive nonmetals; e.g., fluorine (F), chlorine (Cl), bromine (Br), iodine (I).

Ions and the Periodic Table

• Main-group metals tend to lose electrons, forming cations with the same number of electrons as the nearest noble gas.
• Main-group nonmetals tend to gain electrons, forming anions with the same number of electrons as the nearest noble gas.
• Alkali metals (group 1A) tend to lose one electron and form 1+ ions.
• Alkaline earth metals (group 2A) tend to lose two electrons and form 2+ ions.
• Halogens (group 7A) tend to gain one electron and form 1− ions.
• Oxygen family nonmetals (group 6A) tend to gain two electrons and form 2− ions.
• For main-group elements forming cations with predictable charge, the charge equals the group number.
• For main-group elements forming anions with predictable charge, the charge equals the group number minus eight.
• Transition elements may form various ions with different charges.

Atomic Mass

• Atomic mass (or atomic weight/standard atomic weight) is the average mass of the isotopes that compose an element, weighted by their natural abundance.
• Located directly beneath the element’s symbol in the periodic table.
• To calculate atomic mass:
  • Convert percent abundance to decimal form and multiply each by its isotopic mass, then add them.
  • Example: Chlorine
    • 75.77% chlorine-35 (mass 34.97 amu): 0.7577 \times 34.97 \text{ amu} = 26.4968 \text{ amu}
    • 24.23% chlorine-37 (mass 36.97 amu): 0.2423 \times 36.97 \text{ amu} = 8.9578 \text{ amu}
    • Atomic mass Cl = 26.4968 + 8.9578 = 35.45 \text{ amu}
• General equation:
  • \text{Atomic mass} = \sum_{n} (\text{fraction of isotope }n \times \text{mass of isotope }n)

Mass Spectrometry

• Mass spectrometry involves ionizing a sample, accelerating the ions through an electric field, separating them using a magnetic field, and detecting them to determine their mass-to-charge ratio.

Molar Mass

• Chemists need to know the number of atoms in a sample, because chemical processes occur between particles.
• Since it is not practical to count atoms directly, they are counted by weighing.

The Mole

• A mole (mol) is the chemist’s “dozen”.
• It is the measure of material containing 6.02214 \times 10^{23} particles, known as Avogadro’s number.
• 1 \text{ mole} = 6.02214 \times 10^{23} \text{ particles}
• The mole is the number of atoms in exactly 12 grams of pure C-12:
  • 12 \text{ g C} = 1 \text{ mol C atoms} = 6.022 \times 10^{23} \text{ C atoms}

Converting Between Moles and Number of Atoms

• Similar to converting between dozens of eggs and number of eggs.
• Use the conversion factor 1 \text{ mol atoms} = 6.022 \times 10^{23} \text{ atoms}.
• Conversion factors:
  • \frac{1 \text{ mol atoms}}{6.022 \times 10^{23} \text{ atoms}} \text{ or } \frac{6.022 \times 10^{23} \text{ atoms}}{1 \text{ mole atoms}}

Converting Between Mass and Amount

• The mass of 1 mole of atoms of an element is the molar mass.
• An element’s molar mass in grams per mole (g/mol) is numerically equal to the element’s atomic mass in atomic mass units (amu).
• Examples:
  • 26.98 \text{ g aluminum} = 1 \text{ mol aluminum} = 6.022 \times 10^{23} \text{ Al atoms}
  • 12.01 \text{ g carbon} = 1 \text{ mol carbon} = 6.022 \times 10^{23} \text{ C atoms}
  • 4.003 \text{ g helium} = 1 \text{ mol helium} = 6.022 \times 10^{23} \text{ He atoms}

Molar Mass as a Conversion Factor

• Molar mass is the conversion factor between the mass (in grams) of an element and the amount (in moles) of that element.
• For carbon:
  • \frac{12.01 \text{ g C}}{1 \text{ mol C}} \text{ or } \frac{1 \text{ mol C}}{12.01 \text{ g C}}

Conceptual Plan for Counting Atoms by Weighing

• Obtain the mass of the sample.
• Convert to amount in moles using the element’s molar mass.
• Convert to the number of atoms using Avogadro’s number.