Chapter 3

Atoms and Moles

Early Atomic theory

• 4 element theory• Matter is made of atomos• Alchemy• Robert Boyle – proved water could be broken

down; disproved 4 element theory

Laws

• As many new elements were being discovered, scientists realized that there were certain laws that governed how they were made into compounds

Law of definite proportions

• States that a chemical compound always contains the same elements in exactly the same proportions by mass

• In other words, a compound always has the same chemical formula – change the formula, you change the compound

Law of definite proportions (cont)

For example, water is always 88.8% oxygen and 11.2% hydrogen

So the formula for water must be H2O and nothing else

Law of Conservation of Mass

• The mass of the reactants in a reaction equals the mass of the products

• Mass cannot be created or destroyed in ordinary chemical and physical changes.

Dalton’s Atomic Theory• In 1808, John Dalton developed an atomic

theory. • Dalton believed that a few kinds of atoms

made up all matter.• According to Dalton, elements are composed

of only one kind of atom and compounds are made from two or more kinds of atoms.

• Dalton’s model of the atom was a solid sphere

Dalton’s Atomic Theory (cont.)Dalton’s Theory Contains Five Principles

1. All matter is composed of extremely small particles called atoms, which cannot be subdivided, created, or destroyed.

2. Atoms of a given element are identical in their physical and chemical properties.

3. Atoms of different elements differ in their physical and chemical properties.

Dalton’s Atomic Theory (cont.)Dalton’s Theory Contains Five Principles (cont.)

4. Atoms of different elements combine in simple, whole-number ratios to form compounds.

5. In chemical reactions, atoms are combined, separated, or rearranged but never created, destroyed, or changed.

• Data gathered since Dalton’s time shows that the first two principles are not true in all cases.

Subatomic Particles

• Experiments by several scientists in the mid-1800s led to the first change to Dalton’s atomic theory. Scientists discovered that atoms can be broken into pieces after all.

• The smaller parts that make up atoms are called subatomic particles.

• The three subatomic particles that are most important for chemistry are the electron, the proton, and the neutron.

Subatomic Particles, continuedElectrons Were Discovered Using Cathode Rays

• To study current, J. J. Thomson pumped most of the air out of a glass tube. He applied a voltage to two metal plates, called electrodes, which were placed at either end of the tube.

• One electrode, called the anode, was attached to the positive terminal of the voltage source, so it had a positive charge.

• The other electrode, called a cathode, had a negative charge because it was attached to the negative terminal of the voltage source.

CRT

Subatomic Particles, continuedElectrons Were Discovered Using Cathode Rays

• Thomson observed a glowing beam that came out of the cathode and struck the anode and the nearby glass walls of the tube. • He called these rays cathode rays. • The glass tube Thomson used is known as a

cathode-ray tube (CRT). • CRTs are used in television sets, computer monitors,

and radar displays.

Subatomic Particles, continuedAn Electron Has a Negative Charge

• Because the cathode ray came from the negatively charged cathode, Thomson reasoned that the ray was negatively charged.

• Thomson confirmed this prediction by seeing how electric and magnetic fields affected the cathode ray.

• Thomson also observed that when a small paddle wheel was placed in the path of the rays, the wheel would turn.

• This suggested that the cathode rays consisted of tiny particles that were hitting the paddles of the wheel.

Subatomic Particles, continuedAn Electron Has a Negative Charge

• Thomson’s experiments showed that a cathode ray consists of particles that have mass and a negative charge.

• These particles are called electrons. • An electron is a subatomic particle that has a

negative electric charge.• Electrons are negatively charged, but atoms have no

charge.• Atoms contain some positive charges that

balance the negative charges of the electrons.

• Thomson’s model of the atom

Rutherford

• Ernest Rutherford performed the gold foil experiment, which disproved the plum-pudding model of the atom.

• A beam of small, positively charged particles, called alpha particles, was directed at a thin gold foil.

• Rutherford’s team measured the angles at which the particles were deflected from their former straight-line paths as they came out of the foil.

• Rutherford found that most of the alpha particles shot at the foil passed straight through the foil. But very few were deflected, in some cases backward.

Gold Foil Experiment

Gold Foil Experiment (cont)

• Rutherford reasoned that only a very concentrated positive charge in a tiny space within the gold atom could possibly repel the fast-moving, alpha particles enough to reverse the alpha particles’ direction.

• Rutherford also hypothesized that the mass of this positive-charge containing region, called the nucleus, must be larger than the mass of the alpha particle.

• Rutherford argued that the reason most of the alpha particles were undeflected, was that most parts of the atoms in the gold foil were empty space.

Gold Foil Experiment (cont)

Rutherford

• The nucleus is the dense, central portion of the atom.

• The nucleus is made up of protons and neutrons.

• The nucleus has all of the positive charge, nearly all of the mass, but only a very small fraction of the volume of the atom.

Rutherford Model

• Rutherford added a nucleus with protons to the atom

Bohr

• Rutherford did not know why electrons did not get attracted to nucleus

• Was solved by Bohr– Put electrons in orbitals– Said electrons were spinning very fast– Planetary model

Exciting electrons

• Adding energy excites electrons• When an electron becomes excited, it

jumps to a higher energy level• the electron cannot stay in that excited

state forever and so it must fall back to its original energy level

• The energy is then released as light– Law of conservation of energy

Exciting electrons

Atomic Number and Mass Number

• Atomic # = # of protons– Is the identifying characteristic of an atom– In a neutral atom it also = # of electrons

• Mass # = # protons + # neutrons

Isotopes

• Isotopes– An atom that has the same # of protons, but a

different # of neutrons– Example:

• Carbon-12 has 6 protons and 6 neutrons• Carbon-14 has 6 protons and 8 neutrons

Quantum Numbers

• Schrondinger – father of quantum mechanics

• Devised quantum numbers to define the region in which an electron can be found

• There are 4 quantum numbers• No 2 electrons can have the same 4

quantum numbers

Quantum numbers (cont)

1. 1st quantum number = principal quantum number

• Symbol = n• Indicates the energy

level the electron occupies

• n = 1, 2, 3, …

Quantum numbers (cont)

2. 2nd quantum number = angular quantum number

• Symbol = l• Indicates the shape

of the orbital• s = spherical• p = dumbbell• d = double dumbbell• f = star/asterik

Quantum numbers (cont)

3. 3rd quantum number = magnetic quantum number

• Symbol = m• Indicates the orbital’s orientation around the

nucleus• S has 1 • P has 3 • D has 5 • F has 7

Quantum numbers (cont)

4. 4th quantum number = spin quantum number

• Symbol = ms

• Indicates the spin direction of the electron• Since each orbital orientation can only hold

2 electrons, ms can only = +1/2 or -1/2

• Each electron must have opposite spins

Electron configuration

• Pauli exclusion principal– Each orbital can only hold a maximum of 2

electrons

Electron configuration (cont)

• The Aufbau Principle– Electrons fill orbitals that have the lowest

energy first– The smaller the principal quantum number,

the lower the energy

Electron configuration (cont)

• Electron configuration tells the location of each electron– Its energy level– The shape of its orbital– Which orbital it is in– The direction it is spinning

Electron configuration (cont)• Placing the electrons of

an element in their proper location:– Start with the lowest

energy level first (closest to the nucleus) and work your way out

From: http://www.ccs.k12.in.us/chsBS/kons/kons/what_is_organic_chemistry.htm

Electron configuration (cont)

• Fill in electrons in order• Remember that:

– s can only hold 2 (1 orbital x 2 electrons)

– p can hold 6 (3 orbitals x 2 electrons each)

– d can hold 10 (5 orbitals x 2 electrons)

– f can hold 14 (7 orbitals x 2 electrons)

Order of energy levels

Counting Atoms

• Masses of atoms are expressed as atomic mass units (amu)

Moles

• Mole is the number of atoms in exactly 12 grams of carbon-12– This is equivalent to 6.02 x 1023 atoms

• Avogadro’s number

– It is the SI unit for amount of a substance– Molar mass = mass in grams of 1 mole of an

element or compound• This equals the atomic mass of an element

Converting

• Grams/molar mass = moles• Moles x molar mass = grams• Moles x 6.02 x 1023 = atoms• Atoms/ 6.02 x 1023 = moles