ATOMIC STRUCTUREIB Chemistry 2

Robinson High School

Andrea Carver

THE ATOM: IB OBJECTIVES 2.1.1 State the position of protons, neutrons, and

electrons in atoms. 2.1.2 State the relative masses and relative charges

of protons, neutrons, and electrons. 2.1.3 Define the terms mass number (A), atomic

number (Z), and isotopes of an element. 2.1.4 Deduce the symbol for an isotope given its

mass number and atomic number. 2.1.5 Calculate the number of protons, neutrons,

and electrons in atoms and ions from the mass number, atomic number, and charge.

2.1.6 Compare the properties of the isotopes of an element.

2.1.7 Discuss the uses of radioisotopes.

DEVELOPMENT OF ATOMIC THEORY

Democritis- Idea of “atomos.”

John Dalton-Originator of atomic theory.

J.J Thomson-Electrons, Plum-Pudding model.

Ernest Rutherford-Nucleus, Gold Foil Experiment

Niels Bohr-Hydrogen atom model based on emission spectrum.

THE ATOM: SUBATOMIC PARTICLES

Three subatomic particles: Proton-

Positively charged (1+) Mass close to 1 amu Located in the nucleus

Neutron- Neutral/No charge Mass close to 1 amu Located in the nucleus

Electron- Negatively charged (1-) Mass insignificant (0.0005 relative to proton/neutron

mass) Located around the nucleus in “electron cloud”

ATOMIC NUMBER AND MASS NUMBER

The atoms of each element have a characteristic number of protons, represented by the atomic number (Z).

The atomic number also equals the number of electrons present in a neutral atom.

The mass number (A) is the total number of protons and neutrons in the atoms.

Notation:

A

ZX

Mass Number

Atomic Number

Chemical Symbol

ISOTOPES

Isotopes are atoms with identical atomic numbers and different mass numbers.

Isotopes of Carbon:Symbol Number of

ProtonsNumber of Electrons

Number of Neutrons

11C 6 6 5

12C 6 6 6

13C 6 6 7

14C 6 6 8

CALCULATING AVERAGE ATOMIC MASS

Just like carbon, most elements occur in nature as mixtures of isotopes.

The masses of each isotope as well as its relative abundance is taken into account when calculating the average atomic mass (atomic weight) of the element.

The atomic weight is calculated by multiplying the mass of each isotope by its respective percent abundance, and then summing those values.

For example, naturally occurring carbon is composed of 98.93% 12C and 1.07% 13C. Calculate the atomic weight of carbon.

PROPERTIES OF ISOTOPES

A difference in the number of neutrons will have no effect on the atom’s reactivity.

This is because neutrons are not involved in any bonds which the element may form or break in a chemical reaction.

Thus, all isotopes of an element have essentially the same chemical properties.

The difference in mass of isotopes does lead to different physical properties.

A mixture of isotopes may be separated by physical properties such as boiling point or diffusion rate.

RADIOISOTOPES

The stable nuclei of elements fall within the band of stability on a graph of number of protons and number of neutrons.

A stable nucleus requires a balance between protons and neutrons.

Elements with a proton::neutron outside of the band of stability spontaneously emit radioactive particles to gain stability (i.e. move within the band of stability) Radioactive Particles:

Alpha Particle- consists of 2 protons and 2 neutrons Beta Particle- electrons ejected following neutron

decay Gamma Ray- a form of electromagnetic radiation

RADIOISOTOPES: USES

Radioactive isotopes may be used to: Generate energy in nuclear power stations Sterilize surgical instruments in hospitals Preserve food Fight crime Detect cracks in structural materials

Examples: Carbon-14 dating Cobalt-60, radiotherapy Iodine-131, medical tracer

IONS

In an uncharged atom, the atomic number equals the number of protons which equals the number of electrons. (Positive and negative charges are balanced.)

To gain stability, some atoms will gain or lose electrons.

An ion is an atom in which electrons have been lost or gained. These atoms will have an unbalanced charge. Cation- positively charged, has lost electrons Anion- negatively charged, has gained electrons

The ions of an element will have different chemical properties.

THE MASS SPECTROMETER: IB OBJECTIVES

2.2.1 Describe and explain the operation of a mass spectrometer.

2.2.2 Describe how the mass spectrometer may be used to determine relative atomic mass using the 12C scale.

2.2.3 Calculate non-integer relative atomic masses and abundance of isotopes from given data.

THE MASS SPECTROMETER

The most accurate method of measuring atomic and molecular weight is by use of the mass spectrometer.

Relies on the use electric and magnetic fields. Works by five basic operations:

Vaporization-Sample injected as gas. Ionization- Atoms become positively charged via

collision with electrons. Acceleration- Positive ions increase velocity due

to attraction to negative plates. Deflection- Magnetic fields change path of ions. Detection- Positive ions are detected and signal

is recorded.

THE MASS SPECTROMETER

The Carbon-12 atom is used as a standard for comparison for calculating mass of isotopes. Carbon is common, easy to transport and store, and

is solid. Carbon-12 is given a relative atomic mass of 12. Mass spectrum- graphical results of mass

spectrometry X-Axis- mass/charge ratio Y-Axis- % abundance

Practice: The mass spectrum of gallium shows that in a sample of 100 atoms, 60 had a mass of 69, and 40 had a mass of 71. Calculate the average mass of gallium.

ELECTRON ARRANGEMENT: IB OBJECTIVES

2.3.1 Describe the electromagnetic spectrum.

2.3.2 Distinguish between a continuous spectrum and a line spectrum.

2.3.3 Explain how the lines in the emission spectrum of hydrogen are related to electron energy levels.

2.3.4 Deduce the electron arrangement for atoms and ions up to Z=20.

ATOMIC THEORY AND THE ELECTROMAGNETIC SPECTRUM

Flame tests are often used to identify unknown compounds.

This is possible because all elements produce unique emission spectra.

Electrons which have been excited to higher energy levels release excess energy as photons of light as each returns to its ground state.

The amount of energy present in the released photon results in the appearance of a specific color of visible light.

THE ELECTROMAGNETIC SPECTRUM

Visible light is an example of electromagnetic radiation.

Just like any waveform, electromagnetic waves can be described using the terms wavelength, frequency, and amplitude.

THE ELECTROMAGNETIC SPECTRUM

CONTINUOUS SPECTRUM

A source of radiant energy typically emits more than one wavelength of light.

This light can be separated into its component wavelengths.

A prism produces a continuous range of colors from white light. This is called a continuous spectrum.

LINE SPECTRA

Not all radiation sources produce a continuous spectrum.

A line spectrum is a spectrum containing radiation of only specific wavelengths.

The line spectrum of hydrogen initially consisted of four lines: violet-410nm, blue-434nm, blue-green-486nm, and red-656nm.

BOHR MODEL

Bohr proposed the atomic model in which electrons orbit the nucleus much like planets orbit the sun.

Model is based on the following postulates: Only orbits of certain radii corresponding with

specific energies are allowed. An electron in an allowed orbital has a specific

energy and will not emit energy or spiral toward the nucleus.

Energy is emitted or absorbed by an electron as it moves from one allowed energy state to another.

Bohr’s model explains the emission spectrum of hydrogen.

WRITING ELECTRON ARRANGEMENTS

Electrons are added filling energy levels from lowest to highest.

First energy level can hold up to 2 electrons. All subsequent energy levels can hold up to 8

electrons. This method works for elements up to Z=20.

ELECTRON CONFIGURATION:IB OBJECTIVES 12.1.1 Explain how evidence from first ionization

energies across periods accounts for the existence of main energy levels and sublevels in atoms.

12.1.2 Explain how successive ionization energy data is related to the electron configuration of an atom.

12.1.3 State the relative energies of s, p, d, and f orbitals in a single energy level.

12.1.4 State the maximum number of orbitals in a given energy level.

12.1.5 Draw the shape of an s orbital and the shapes of px, py, and pz orbitals.

12.1.6 Apply the Aufbau principle, Hund’s rule, and Pauli exclusion principle to write electron configurations for atoms and ions up to Z=54.

QUANTUM THEORY

Modern atomic theory relies on quantum mechanics to describe interactions and locations of electrons in space.

The probably location of an electron within an atom can be described by four quantum numbers. The Principle Quantum Number (n) The Angular Quantum Number (l) The Magnetic Quantum Number (m) The Spin Quantum Number

SUB-LEVELS

Each energy level contains as main sublevels as its n, so the first principal energy level contains one sublevel, the second contains two, and so on.

The letters s, p, d, and f are used to distinguish the sublevels within energy levels.

Each sublevel can contain a characteristic number of electrons: S-2 P-6 D-10 F-14

THE UNCERTAINTY PRINCIPLE

Heisenberg’s Uncertainty Principle states that it is impossible to know the location of an electron at any moment in time.

This is because electrons are moving and any attempt to measure an electrons location or movement would change its location or movement.

We are only able to predict an area within which an electron is likely to be.

We refer to this area as an atomic orbital.

S ATOMIC ORBITAL

P ATOMIC ORBITALS

D AND F SUB-LEVELS

ELECTRON CONFIGURATION RULES

Pauli Exclusion Principle- No two electrons in an atom can have the same four quantum numbers.

Aufbau Principle- An electron occupies the lowest energy level that is able to receive it.

Hund’s Rule- For degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin is maximized. (Seats on a bus.)

WRITING ELECTRON CONFIGURATIONS

Electron configurations can be written using orbital notation.

Each orbital is represented by a ____, and each electron will be represent by an arrow.

Electrons are distributed according to the Aufbau Principle, Pauli Exclusion Principle, and Hund’s Rule.

WRITING ELECTRON CONFIGURATIONS

Tedious lines and arrows are eliminated by using electron configuration notation.

Sublevels are indicated by letter, s,p,d, or f. The number of electrons in a sublevel is

indicated by use of a subscript. The principle energy level is indicated by a

coefficient. It is possible to write electron configurations

simply by using an elements location on the periodic table.