Atomic Structure

1. Atomic Structure IB Chemistry 2 Robinson High School Andrea Carver

2. 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.

3. 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.

4. 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”

5. 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: Mass Number A X Chemical Symbol Z Atomic Number

6. Isotopes • Isotopes are atoms with identical atomic numbers and different mass numbers. • Isotopes of Carbon:

7. 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.

8. 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.

9. 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

10. 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

11. 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.

12. 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.

13. 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.

14. 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.

15. 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.

16. 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.

17. 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.

18. The Electromagnetic Spectrum

19. 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.

20. 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.

21. 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.

22. 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.

23. 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 pzorbitals. • 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.

24. 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

25. 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

26. 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.

27. S Atomic Orbital

28. P Atomic Orbitals

29. D and F Sub-Levels

30. 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.)

31. 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.

32. 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.