SC3. Students will use the modern atomic theory to explain the characteristics of atoms.a. Discriminate between the relative size, charge, and position of protons, neutrons, and electrons in the atom.b. Explain the relationship of the proton number to the element’s identity.c. Explain the relationship of isotopes to the relative abundance of atoms of a particular element.

The Development of the Atomic Theory

THE INDIVISIBLE ATOMGreek - Democritus (450 BC)

- Atom was indivisible- Theorized the existence of the atom - Also, theorized that there were just four 'elements' - fire, water, air, earth

John Dalton (1803) - Atom was indivisible- All elements are composed of atoms- The same atoms for one element are exactly alike- Atoms are neither created or destroyed in a chemical reaction- In a chemical reaction, atoms are separated, combined, or rearranged- Law of Definite Proportions – atoms combine in definite whole number ratios by mass to

form compounds. The ratio of one element to another in a compound is always definite. Example:

- Law of Multiple Proportions – atoms react as whole parts and cannot be divided into smaller parts in a chemical reaction. Example:

-

THE DIVISIBLE ATOM

J. J. Thomson (late 1890's)- discovered the electron using the cathode ray tube - determined that the electron was smaller than a hydrogen atom. This was a shocking

discovery. Many thought Dalton was wrong.- Knew the atom was neutral and the electron was negative, so there must be positive material

with a lot more mass. Said the atom was a positive pudding-like material throughout which negatively charged electrons were scattered - Plum Pudding or Chocolate Chip Cookie Model ( Page 94, figure 4-9).

VIDEODISC – Rutherford’s Experiment1. What is an alpha particle?

2. What did Rutherford expect to see? What did he actually see?

3. What is Rutherford’s model of the atom?

Ernest Rutherford (1909) (1919)- Did a famous gold foil experiment (the alpha scattering experiment)- Calculated that the atom was mostly empty space through which electrons move. - Concluded that the atom has a small, dense, positively charged, centrally-located nucleus

surrounded by negatively charged electrons- By 1919 he had refined the concept of the nucleus. Called the positive particles protons.

Rutherford and James Chadwick (1932)- showed the nucleus also had a neutron.- The neutron was basically equal in mass to the proton but had no electrical charge.

Neils Bohr (1914)- Electrons moved around the nucleus in definite orbits or energy levels; called the planetary

atomic model.

The atom at this point: - is spherically shaped with a dense, centrally located, positively charged nucleus surrounded

by one or more negatively charged electrons in an electron cloud. - Most of the atom consists of fast moving electrons traveling through empty space

surrounding the nucleus. - Electrons are held within the atom by an attraction to the positive nucleus. - The nucleus has neutral neutrons and positive protons and 99.9% of the mass. - Since atoms are electrically neutral, the number of protons must equal the number of

electrons.

SUBATOMIC PARTICLES

Nucleus - center of the atom- 99.9% of the mass of the atom- contains two particles

Proton - positive charge- all protons are identical, regardless of the element- 1 amu is its mass- the number of protons determines which element you have. If the number of protons

changes, so does the element. Why? Because the number of protons equals the Atomic Number which is unique for every element

Neutron - neutral (no charge)- all neutrons are identical, regardless of the element - 1 amu is its mass

Electron Cloud - space where the electron is most likely to be- electrons do not move in fixed paths around the nucleus- the location of the electron depends on how much energy the electron has- Electron Levels: how electrons are arranged; Each level can hold only a certain

amount of electrons. There are seven levels. Level 1 holds 2 electrons; Level 2 and 3 holds 8 electrons each. When level one fills up, the electrons go to the second level and then the third, etc. (Kind of like filling up a shelf with books. When the first shelf is filled, you move to the second shelf, etc.)

*- Chemical properties of the different elements depends on the number of electrons at the various levels

Electron - located outside of the nucleus in the electron cloud- negative charge- 1/1840 amu

Henry Moseley (1912)- Discovered that atoms of each element contain a unique positive charge. - The number of protons in an atom identifies it.- ATOMIC NUMBER = number of protons (and number of electrons). It is unique for each

elementKEY FOR PERIODIC TABLE

1

HHydrogen

1.01

Atomic Number and Mass NumberYou can figure a lot about an atom by its ATOMIC NUMBER and its MASS NUMBER. The atomic number of an element NEVER changes; the number of protons is always the same. The number of neutrons and electrons can change.

ATOMIC NUMBER = Number of Protons

MASS NUMBER = Number of protons and neutrons (called nucleons)

NUMBER OF PROTONS = Number of electrons in a neutral atom (meaning the number of negative electrons equals the number of positive protons)

NUMBER OF NEUTRONS = Mass Number minus Atomic Number

ISOTOPESIsotopes are atoms of the same element that have a different number of neutrons. They differ in

mass, but the atom’s chemical behavior are the same. To identify an isotope, you write the element’s name and follow it with its mass number. Examples: Carbon-12, Uranium-238. It can also be identified by writing

126 C 23

8 U

Isotopes or different elements?1. Element D 6 p+ and 7 n Element F 7 p+ and 7 n 2. Element J 27 p+ and 32 n Element L 27 p+ and 33 n 3. Element X 17 p+ and 18 n Element Y 18 p+ and 17 n 4. Element T Z = 20 and A = 40 Element Z Z = 20 and A = 41 5. Element P Z = 92 and A = 238 Element S 92 p+ and 143 n

PRACTICEThe atomic mass of iridium-191 is 191.0 amu and iridium-193 is 193.0 amu. The percentage abundance for each is 37.58% (iridium-191) and 62.62% (iridium-193). Calculate the average atomic mass of iridium.

The mass of individual atoms is too small to measure so the mass of an atom is compared to a standard, Carbon-12 – Carbon-12 = 12 amu. AMU stands for atomic mass unit.

The ATOMIC MASS is the weighted average of the isotopes of an element and can be calculated if you know the isotope’s mass numbers and the percentage abundance of each.

SC1. Students will analyze the nature of matter and its classifications.a. Relate the role of nuclear fusion in producing essentially all elements heavier than hydrogen.

NUCLEAR CHANGE

Review: Remember that a chemical reaction involves a change of one or more substances into new substances. Chemical reactions involve only an atom’s electrons. The nucleus remains unchanged.

Characteristics of Chemical and Nuclear Reactions

Chemical Reactions1. Occur when bonds are broken and formed2. Atoms remain unchanged, although they

may be rearranged3. Involve only valence electrons4. Associated with small energy changes5. Reaction rate is influenced by pressure,

temperature, concentration, and catalysts.

Nuclear Reactions1. Occur when nuclei emit particles and/or rays2. Atoms are often converted into atoms of

another element3. May involve protons, neutrons, and

electrons4. Associated with large energy changes5. Reaction rate is not normally influenced by

pressure, temperature, and catalysts.

HISTORY1896 Henri Becquerel - discovered mysterious rays coming from uranium. Called it RADIATION. 1898 Marie and Pierre Curie discovered these rays in Radium and Polonium

Nuclear reactions involve a change in the nucleus. Radioactivity is the spontaneous emission of radiation from an element. They do this because their nuclei are unstable. The rays and particles emitted are called RADIATION.

Unstable nuclei lose energy by emitting radiation in a spontaneous (does not require energy) process called radioactive decay. The atom undergoes decay until it becomes stable. By emitting radiation, atoms of one element change into atoms of another element.

Elements with an atomic number greater than 83 are radioactive. Elements with #1-82 have isotopes that may be radioactive.

TYPES OF RADIATION1. Alpha radiation, Made of alpha particles. Is composed of two protons and two neutrons. Has a 2+ charge and a mass of 4 amu. Has the least amount of energy of any of the radiation – is stopped by paper. A new element is created when alpha decay happens. The mass number and the atomic number change.Alpha decay: 22

6 Ra 222 Rn + \s\up 8(4 ) He Radium-226 decays to radon-222 and an alpha

particle

2. Beta radiation, Made of fast moving electrons. Has a –1 charge and a mass of 1/1840. An electron is emitted during beta decay because it has been removed from a neutron, leaving behind a proton. Is stopped by foil and has more energy than alpha radiation. Beta decay: 22

11Na 2212Mg + \s\up 8(0 ) Sodium-22 decays into magnesium-22 and a beta

particle

3. Gamma radiation, High energy radiation that possesses no mass. Has no charge. Usually accompanies alpha or beta radiation. Is slowed down by lead or concrete. Accounts for most of the energy lost in a radioactive decay process.

Gamma decay: 238 U 23

4 Th + \s\up 8(4 )He + \s\up 8(0 ) Uranium-238 decays into thorium-234 and an alpha

particle and 2 gamma rays of different frequencies

NUCLEAR STABILITY

The ratio of neutrons and protons determines nuclear stability. Too few or too many neutrons make the nucleus unstable. It needs to lose energy to become stable. So it decays. Few radioactive elements are found in nature – they have already decayed into stable atoms.

The STRONG NUCLEAR FORCE acts only on subatomic particles that are extremely close together. It overcomes the electrostatic repulsion between positive protons. Neutrons add an attractive force within the nucleus because they are not positive or negative. They are also subjected to the strong nuclear force. They help hold the nucleus together.

After radioactive decay, the new atom is now positioned more closely, if not within, the band of stability.

THREE TYPES OF NUCLEAR REACTIONS1. SPONTANEOUS - a single nucleus releases energy and particles of matter2. FISSION - nucleus is split in two3. FUSION - two nuclei join together – this is how all elements are made.

*ALWAYS involve a change in the nucleus and a release of energy

RADIOACTIVE DECAYNaturally occurring (radioactive isotopes) radioisotopes are not uncommon on earth.

Radioactive decay rates are measured in half-lives. A HALF-LIFE is the time it takes for one half of the radioactive sample to decay. Different isotopes have different half-lives. Examples include uranium-238 (4.46 x 109 years), carbon-14 (5730 years), radon-222 (3.8 days) and polonium-214 (163.7 microseconds).

If you were given a 10.0g sample of carbon-14, how much would be stable after 3 half-lives?

radioactivestable

original sample 1st half-life 2nd half-life 3rd half-life

5.0g stable 7.5g stable 8.75g stable

WRITING AND BALANCING NUCLEAR EQUATIONS PRACTICERemember when balancing these equations that mass number and atomic number are conserved.

1. Write the nuclear equation for the alpha decay of polonium-208.

2. Write the nuclear equation for the beta decay of copper-66.

SC3. Students will use the modern atomic theory to explain the characteristics of atoms.b. Relate light emission and the movement of electrons to element identification.

WAVE NATURE OF LIGHT

Rutherford’s model of the atom did not explain how the electrons of an atom are arranged in the space around the nucleus. His model did not take into account the chemical behavior among various elements. In the early 1900’s it was observed that certain elements emitted visible light when heated in a flame. Analysis of the emitted light revealed that an element’s chemical behavior was related to the electron arrangement in the atom.

Electromagnetic radiation is a form of energy that exhibits wavelike behavior as it travels through space. Radiant energy, which includes visible light, is shown in the electromagnetic spectrum. Lowest HighestFrequency frequency104 1022

Radio Microwave Radar Infrared Visible Light Ultraviolet X-ray Gamma

longest ROYGBIV shortestwavelength increasing energy wavelength104 10-12

Wave Characteristics:1. wavelength, measured from crest to crest, measured in m, cm, nm2. frequency, measured by number of waves per second, hertz, Hz, waves/s4. speed, c 3.00 x 108 m/s in a vacuum. Speed is always the same. Waves can

have different wavelength and frequency. Wavelength and frequency are inversely related. speed = wavelength x frequency c =

White light can be refracted into its component colors and each color has its own wavelength.

PRACTICE:1. Calculate the speed of a wave whose wavelength is 1.5 meters and whose frequency is 280 hertz.

2. Find the wavelength of a wave whose speed is 5.0 m/s and whose frequency is 2.5 hertz.

3. The speed of light is 3.00 x 108 m/s. Red light has a wavelength of 7.0 x 10-7m. What is its frequency?

PARTICLE NATURE OF LIGHT

Wave model cannot explain why heated objects emit only certain frequencies of light at a given temperature or why some metals emit electrons when colored light of a specific frequency shines on them.1900 Max Planck – Matter can gain or lose electrons only in small specific amounts called quanta. A

quantum is the minimum amount of energy that can be gained or lost. Glowing objects emit light, which is a form of energy.

Energyquantum = h h = Planck’s constant = 6.626 x 10-34 J·s

The energy of the radiation increases as the frequency increases. This explains why ultraviolet light has more energy than violet light. The particle of light is called a photon.

PRACTICE:1. Calculate the energy of a gamma ray photon whose frequency is 5.02 x 1020 Hz.

2. What is the difference in energy between a photon of violet light with a frequency of 6.8 x 1014 Hz and a photon of red light with a frequency of 4.3 x 1014 Hz?

ATOMIC EMISSION SPECTRA

A neon light works because electricity is passed through the neon gas in the tube. The gas absorbs energy and becomes excited. Excited and unstable atoms then release energy by emitting light.

Atomic emission spectrum is a set of frequencies of electromagnetic waves emitted by atoms of the element. They are usually distinct color lines. Each element’s atomic emission spectrum is unique, can be used to identify the element, and can be used to determine if the element is part of an unknown compound.

QUANTUM THEORY

In 1913, Neils Bohr came up with the quantum model of the hydrogen atom. He also correctly predicted the frequencies of the spectral lines in the hydrogen atomic emission spectrum.

The hydrogen atom has only certain allowable energy states. The lowest is called the GROUND STATE. When atoms gain energy, they are said to be in an EXCITED STATE. Hydrogen has one electron, but can have different excited states.

Bohr said that the electron moves around the nucleus in only certain allowable circular orbits. The smaller the orbit, the lower the atom’s energy state or energy level. Each orbit is assigned a quantum number, n.

Bohr said that hydrogen’s electron in the ground state did not radiate energy. When excited, the electron moves up to another energy level. Only certain atomic energies are possible and so only certain frequencies of electromagnetic radiation can be emitted.

Werner Heisenberg concluded that it is impossible to make any measurement on an object without disturbing the object. It is fundamentally impossible to know precisely both the velocity and position of a particle at the same time.

Erwin Schödinger, in 1926, developed an atomic model in which the electrons are treated as waves. This model is called the quantum mechanical or wave mechanical model. This model does not describe the path of the electron around the nucleus. The three dimensional region around the nucleus called an atomic orbital describes the electron’s probable location. The electron cloud has no definite boundary – it is arbitrarily drawn at 90

QUANTUM MECHANICAL MODEL OF THE ATOM1.2. The atom has a dense, centrally located, positively charged nucleus full of protons and neutrons

surrounded by mostly empty space called an electron cloud where the electrons are. (Rutherford)3. The energy of electrons is quantized (has only specific amounts of energy). (Bohr)4. Electrons exhibit both wave and particle behaviors. (de Broglie)5. The absolute location of an electron is impossible to determine – its location and velocity cannot be

determined at the same time. (Heisenberg)6. The electrons travel in orbitals that have characteristic sizes, shapes, and energies, but do not describe

how the electrons move. (Schödinger)

QUANTUM NUMBERSThese numbers describe the most probable location of an electron in an atom.

1. Principle quantum number, nIs the energy level number. Gives information about the relative size and energy of the atomic orbitals. It can have values of 1, 2, 3, 4…. The greatest number of electrons possible in any one level is 2n2.Example: The maximum number of electrons that can occupy the first level is 2(1)2 = 2; the fourth level is 2(4)2

= 32.

2. Angular Momentum Quantum Number, lIs the energy sublevel number. It gives information as to the shape of the orbitals. The first four levels are s, p, d, and f.Example: The first energy level has only an s sublevel. The second energy level has an s and p sublevel. This

third energy level has s, p, and d sublevel.

VIDEODISC:1. What is the Heisenberg uncertainty principle?

2. How is the principle quantum number related to the size and energy of an atomic orbital?

3. What is the electron configuration of iron?

3. Magnetic Quantum Number, mgives information about the orientation in space of an orbital. The s sublevel has one orbital, the p sublevel has three orbitals, the d sublevel has 5 orbitals, and the f sublevel has 7 orbitals. This number determines which p, d, or f orbital the electrons are in.

4. Electron Spin Quantum number, sindicates the direction of the electron spin. Spin is either clockwise or counterclockwise and is designated with a +1/2 or a –1/2 .

ELECTRON CONFIGURATION

Electron configuration tells us how electrons are distributed among the various atomic orbitals. It is the arrangement with the lowest possible energy. It is a simple way of keeping track of electrons. Electrons fill the levels according to a set of rules:

1. Aufbau PrincipleEach electron occupies the lowest energy orbital available. Electrons are added from the ground state up. Electrons fill in increasing energy order.

2. Pauli Exclusion PrincipleEach orbital can hold a maximum of only two electrons – one spinning clockwise and one spinning counter clockwise.

3. Hund’s RuleThe most stable arrangement of electrons in orbitals is to fill singly and then go back and double up.

Examples:1 Hydrogen

2 Helium

3 Lithium

4 Beryllium

5 Boron

6 Carbon

7 Nitrogen

8 Oxygen

9 Fluorine

10 Neon

11 Sodium

12 Magnesium

13 Aluminum

The order that the levels and sublevels fill is based on energy. They fill in the following order:1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p

There are two diagrams to help you remember the filling order: the diagonal rule or the arrow diagram. Can also use the periodic table.

ODDBALL ELECTRON CONFIGURATIONFOR ODDBALLS AND IONS

The transition elements do not always follow the rules for electron configuration. Notice the electron configurations for elements 21-30. The ones in bold are oddballs:

Scandium 1s22s22p63s23p64s23d1

Titanium 1s22s22p63s23p64s23d2

Vanadium 1s22s22p63s23p64s23d3

Chromium 1s22s22p63s23p64s13d5

Manganese 1s22s22p63s23p64s23d5

Iron 1s22s22p63s23p64s23d6

Cobalt 1s22s22p63s23p64s23d7

Nickel 1s22s22p63s23p64s23d8

Copper 1s22s22p63s23p64s13d10

Zinc 1s22s22p63s23p64s23d10

The chromium and copper do not follow the rules. The “d” sublevel holds ten and is more stable when is it half full then the “s” when it is full. So it steals an electron from “s” orbital to be more stable. The same thing happens as it nears ten electrons to be full.

Now, you write the electron configurations for some elements listed below.

42-Molybdenum _______________________________________________________________

43-Technetium _______________________________________________________________

44-Ruthenium _______________________________________________________________

45-Rhodium _______________________________________________________________

46-Palladium _______________________________________________________________

47-Silver _______________________________________________________________

48-Cadmium _______________________________________________________________

Name the transition elements that will be oddballs that you must remember:

ELECTRON CONFIGURATION FOR IONS

An ion is an atom that has a charge. It is not neutral and its number of protons does not equal its electrons. If it has lost electrons, it is positive and called a CATION. If it has gained electrons, it is negative and called an ANION. For atoms that occur as ions, the electron configuration can also be written.

1. First, write the regular electron configuration for the element.2. If the ion is positive, take away electrons. If the ion is negative, add electrons to the highest partially filled

energy level.

EXAMPLES:Na 1s22s22p63s1 Al 1s22s22p63s23p1

Na+1 1s22s22p6 Al+3 1s22s22p6

F 1s22s22p5 H 1s1

F-1 1s22s22p6 H-1 1s2

PRACTICE:

Magnesium, Mg Magnesium ion, Mg+2

Oxygen, O Oxygen ion, O-2

Lithium, Li Lithium ion, Li+1

Iron, Fe Iron ion, Fe+3

Atoms that have absorbed energy and are in an excited state have electrons that have moved to a shell level that is higher than what is normal. The electrons that move come from the valence electrons (outermost electrons).

ORBITAL DIAGRAMS

This is another means of symbolizing where electrons are in energy levels and sublevels. Involves three basis symbols:

Unoccupied orbital orbital with one electron orbital with two electrons

Can also be drawn with circles or lines instead of squares.

_____

ELECTRON CONFIGURATION SHORTHAND

Electron configurations can be written in a shorthand form. You can take the noble gas (ONLY) that occurs before the element in question and then tack on the remaining configuration.

Example: Sodium 1s22s22p63s1 Shorthand: [Ne] 3s1

Copper 1s22s22p63s23p64s13d10 Shorthand: [Ar] 4s13d10

VALENCE ELECTRONS

Valence electrons are the electrons in the outermost energy levels. Using the two examples from above, sodium and copper, the number of valence electrons for each are:

Sodium 1s22s22p63s1 Shorthand: [Ne] 3s1 valence electrons: 1Copper 1s22s22p63s23p64s13d10 Shorthand: [Ar] 4s13d10 valence electrons: 1

Other examples:Magnesium 1s22s22p63s2 valence electrons: 2Chlorine 1s22s22p63s23p5 valence electrons: 7

You can also use the periodic table itself to determine the valence electrons for any atom.

ELECTRON DOT DIAGRAMS

Usually we are only concerned with the electrons in the outermost energy level. Those are the electrons involved in chemical reactions – the VALENCE electrons. We can symbolize these electrons using the Lewis Electron Dot Diagram. To write an electron dot diagram, follow these steps:

1. Write the symbol for the element (symbolizes the nucleus and all inner energy levels)2. Write the electron configuration for the element. Select the electrons in the outermost energy level.

Use the HIGHEST principle quantum number regardless of highest energy.3. Each side of the symbol represents an orbital. Draw the dots on the appropriate sides to represent the

electrons in that orbital. It is important to remember which electrons are paired and which are not.

s s

p pp p

p p

Examples:Sodium 1s22s22p63s1 valence electrons: 1 Na

Copper 1s22s22p63s23p64s13d10 valence electrons: 1 Cu

Magnesium 1s22s22p63s2 valence electrons: 2 Mg

Chlorine 1s22s22p63s23p5 valence electrons: 7 Cl

X