1. CHAPTER 5 Electrons in Atoms

2. Development of Atomic Models Dalton – Remember atomic theory? – Atom considered indivisible Thomson – “plum pudding atom” – Electrons stuck into a mass of protons – Did not take into account number of protons, neutrons, and electrons, arrangement within atom, or how ions are formed

3. Development of Atomic Models Rutherford – Nuclear atom in which electrons surround dense positively charged nucleus – Rest of atom empty space – What prevents the electrons from falling into the nucleus?

4. Development of Atomic Models Bohr – “planetary” model of the atom – Electrons in a particular path have a fixed energy; they do not lose energy and fall into the nucleus – Energy level: region around the nucleus where the electron is likely to be moving – Quantum: amount of energy required to move an electron from one energy level to the next higher one.

5. Development of Atomic Models

6. Bright line spectra: when connected to an energy source, gases emit distinct lines of color which can be used to identify elements Hydrogen: H 2 always had the same 4 colors at the same 4 frequencies – electrons can only exist at certain energy levels

7. Development of Atomic Models 1. When “excited”, electrons absorb energy and “jump” to a higher energy level ( quantum ) 2. Excited atoms are unstable and lose energy only in certain amounts 3. Light waves of distinct frequencies are emitted when electrons return to their “normal” energy level

8. Properties of Waves Waves possess energy – Energy is directly proportional to frequency E=hν Energy in joules=(6.63 x 10 -34 J*s)x(frequency in Hz) – Frequency and wavelength are indirectly proportional c=λν 3.0 x 10 8 m/s=(wavelength in m)x(frequency in Hz)

9. Determining Properties of Waves 1. A wave has a frequency of 5.0x10 -2 Hz, what energy does it possess? 2. If a wave has a 3.4x10 -4 m wavelength, what is its frequency? 3. What is the frequency of a wave with 2.93x10 -12 J of energy? 4. What is the energy possessed by a wave with a wavelength of 0.025 mm?

10. Quantum Mechanical Model Quantum mechanical model comes from the mathematical solution to the Schrödinger equation Does not define exact path an electron takes around the nucleus, but does restrict energy of electrons Estimates probability of finding an electron in a certain position

11. Quantum Mechanical Model Principal quantum numbers (n) designate energy levels of electrons Electrons occupy energy sublevels ; number of sublevels = the principal quantum number Atomic orbitals are the cloud shapes where electrons can be found; denoted by letters (s, p, d, and f) Maximum number of electrons that occupy an energy level = 2n 2

12. Electron Configurations The ways in which electrons are arranged around the nuclei of atoms are called electron configurations Aufbau principle: electrons enter orbitals of lowest energy first Pauli exclusion principle: an atomic orbital may describe at most two electrons Hund’s rule: when electrons occupy orbitals of equal energy, one electron enters each orbital until all the orbitals contain one electron with parallel spins

13. Electron Configurations 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f

14. Electron Configurations Write electron configurations for these atoms: phosphorus and nickel phosphorus : 1s 2 2s 2 2p 6 3s 2 3p 3 3p ↑ ↑ ↑ 3s ↑↓ 2p ↓↑ ↓↑ ↓↑ 2s ↑↓ 1s ↓↑

15. Electron Configurations Nickel: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 8 3d ↑↓ ↑↓ ↑↓ ↑ ↑ 4s ↑↓ 3p ↑↓ ↑↓ ↑↓ 3s ↑↓ 2p ↑↓ ↑↓ ↑↓ 2s ↑↓ 1s ↑↓

16. CHAPTER 6 Chemical Periodicity

17. Development of the Periodic Table Scientists needed a way to organize and refer to the growing number of known elements Categorize the elements according to similarities in their physical and chemical properties Relationship between atomic structure and properties of elements

18. Development of the Periodic Table Mendeleev listed the elements in vertical columns in order of increasing atomic mass Regular recurrence of physical and chemical properties Blank spaces left where unknown elements belonged, but able to predict their properties

19. Development of the Periodic Table Moseley determined atomic number of the atoms of the elements Arranged the elements by order of atomic number

20. Modern Periodic Table The horizontal rows of the periodic table are called periods (organized by increasing atomic number) The vertical columns are called groups or families (organized by chemical properties) Periodic Law: when the elements are arranged in order of increasing atomic number, there is a periodic pattern in their physical and chemical properties

21. Electron Configurations and Periodicity Electron plays the greatest part in determining the physical and chemical properties of an element Elements can be classified into four different categories based on their electron configurations

22. Electron Configurations and Periodicity 1. The noble gases are elements in which the outermost s and p sublevels are filled (group 0) – Called inert gases because they do not partake in chemical reactions – Valence: outermost s and p orbitals – Octet: 8 electrons in valence shell

23. Electron Configurations and Periodicity 2. The representative elements are elements whose outermost s or p sublevels are only partially filled (group A) – Alkali metals: group 1A elements – Alkaline earth metals: groups 2A elements – Halogens: nonmetallic elements of group 7A – Group number = number of valence electrons

24. Electron Configurations and Periodicity 3. The transition metals are elements whose outermost s sublevel and the nearby d sublevel contain electrons (group B) 4. The inner transition metals are elements whose outermost s sublevel and the nearby f sublevel contain electrons

25. Electron Configurations and Periodicity S block: groups 1A and 2A and He P block: groups 3A, 4A, 5A, 6A, 7A, and 0 except He D block: transition metals F block: inner transition metals Period number corresponds to the principal energy level

26. Periodic Trends in Atomic Size The radius of atoms: atomic radii Atomic size increases as you move down a group of the periodic table and decreases as you move form left to right across a period The shielding effect of these electrons on the nucleus is constant within a period

27. Periodic Trends in Ionization Energy The energy that is required to overcome the attraction of the nuclear charge and remove an electron to create a positive atom is the ionization energy First ionization energy decreases as you move down a group of the periodic table and increases as you move form left to right across a period

28. Periodic Trends in Ionic Size Ionic radii is the size of an ion compared to same neutral atom – Cations (positive ions) are smaller because of the reduced energy level with the same number of protons – Anions (negative ions) are larger because of a reduced effective nuclear charge on outer electrons

29. Periodic Trends in Electronegativity The electronegativity of an element is the ability of an atom to attract electrons within a bond Electronegativity increases as you go across a period form left to right and decreases as you move down a group