Unit 1 Atomic Structure and Nuclear Chemistry

1. Unit 1 Atomic Structure and Nuclear Chemistry

2. 4.1 - The Elements • A. An element is: • B. Elements can exist as pure substances or as parts of compounds. • Examples: • C. Representative elements (main group) • 4.2 – Symbols for the Elements • Chemical symbols are: • Be forewarned, the one or two letter chemical symbol for many elements is not the same as the first one or two letters in the element name. • Examples:

3. 1. Gold Foil Experiment (Rutherford)

4. 4.6 – Introduction to the Modern Concept of Atomic Structure A. From Rutherford’s gold foil experiment we learned that. 1. Atoms have a central positive nuclear charge 2. Atoms are 99.99% empty space. B. Subatomic Particles

5. E. A “Box” in the Periodic Table:

6. Mass Number = Number of Protons + Number of Neutrons Also, for a neutrally charged atom, Atomic Number = Number of Protons = Number of Electrons *the atomic number determines the identity of an element

7. 4.7 – Isotopes • A. Atoms of the same element always have the same number of ___________________. • *the number of protons an atom of a certain element contains is given by its _____________________________________. • Isotopes are: • The sum of the number of protons and the number of neutrons contained in an atom is know as the atom’s ________________. • D. Isotope Notation: • 1. Isotopes are often symbolized by an element symbol with superscripts and subscripts denoting the isotope’s mass number and atomic number. • Examples: • 2. Isotopes can also be written as the element’s name, a dash, and the mass number of the isotope. • Examples:

8. Use the equations above and your periodic table to fill in the missing values in the table below:

9. 8.1 Weighted Average Atomic Mass A. What is a weighted average? Why do we need one? An element can exist in a number of forms, called isotopes. Isotopes are forms of the same atom that vary in mass.  For example, there are two different types (isotopes) of copper atoms. One type of copper atoms weighs in at 62.93 amu, the other has a mass of 64.94 amu. The lighter isotope is more common with 69.09% of the naturally occurring copper having a mass of 62.93 amu per atom. The remainder of the atoms, 30.91 %, have a mass of 64.94 amu. To find the Average Atomic Mass of an atom, we take into account all of the isotopes that exist and the percentage of each type. The calculation of the average atomic mass is weighted average.

10. B. How to calculate a weighted average atomic mass Take the sum of the products of each isotope’s mass and its corresponding relative abundance as a decimal (take percent and move decimal 2 places left). Example 1

11. Example 2

12. 4.10 Ions • A neutrally charged atom has a(n) _______________________ of protons and electrons. • A neutrally charged atom can become an ion if ___________________________________________________________________________________________ • C. An ion is __________________________________________ • 1. A positively charged ion is known as a ____________________. • 2. A negatively charged ion is an ________________________. • D. Ion charges and the periodic table. • 1. Elements have the tendency to gain or lose electrons to have the same number of electrons as the closest noble gas. • Examples: • 2. You can use the column that an element is in to predict charges (doesn’t work for transition metals) • Examples:

13. 11.1 Rutherford’s Atom A. His model leaves many questions about electrons unanswered: 1. 2. 3.

14. Light and Spectra • Every element emits light when its given energy • Atomic Emission Spectrum = passing the light emitted by an element through a prism

15. How Light Relates to Electron Location • Bohr observed that only certain colors were given off • Therefore the electron could only orbit at certain distances from the nucleus

16. B. Types of EM Radiation and wavelength.

17. 11.3 Emission of Energy by Atoms • Electrons surrounding an atom can absorb a discrete packet of energy called a ___________________________ to become “excited”. • B. When excited electrons lose that extra energy they fall back into their • _______________________. The release of energy by the electron results in emission of a photon of a certain wavelength. Each element has its own unique spectrum of wavelengths that are released. • Examples:

18. Bohr Model

19. 11.5 The Bohr Model of the Atom A. Bohr’s model of the atom included the following main points. 1. Central nucleus made up of ______________ and _______________. 2. Electrons were restricted to circular orbits. • 11.6 The Wave Mechanical Model of the Atom (Quantum Mechanics Model) • This model of the atom is our most current and up to date model. • B. The biggest difference between this model and Bohr’s model is “Orbits vs. Orbitals”. • 1. An Orbit is – • 2. An Orbital is –

20. 11.7 The Hydrogen Orbitals A. Within each principal energy level there can be one or more orbitals. 1. “s” orbitals: 2. “p” orbitals: 3. “d” orbitals: B. Each principle energy level is a little larger and further away from the nucleus than the last and contains more orbitals than the last.

21. Electron Configurations • In the atom, electrons arrange themselves around the nucleus in the most stable way possible • Called electron configurations • Three rules that explain electron configurations: • Aufbau principle • Pauli exclusion principle • Hund’s rule

22. Electron Configurations • Aufbau Principle: • Electrons enter orbitals of lowest energy first • s is ALWAYS the lowest energy level within a principle energy level

23. Electron Configurations • Pauli exclusion principle: • An atomic orbital may have AT MOST two electrons • s = ____ orbital or ___ e- max • p = ____ orbitals or ___ e- max • d = ____ orbitals or ___ e- max • f = ____ orbitals or ___ e- max • Each box ( ) denotes an orbital • For an orbital to be “filled” with electrons, they must have opposite spins

24. Electron Configurations • 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” • Example: Oxygen (8 electrons)

25. Let’s apply the 3 rules with examples of Electron Configurations

28. Electron Arrangements in the First 18 Atoms on the Periodic Table (11.9) • A. Electron Arrangement = Electron Configuration • B. Box Diagram • 1. Principal Energy Level • 2. Type of orbital • 3. Valence Electrons • 4. Core Electrons

29. Electron Configurations and the Periodic Table (11.10) • Exceptions: Cr, Mo, Cu, Ag, Au • Transition Metals (one energy level below) • Lanthanide and Actinide Series (two energy levels below) • Note the following points: • Groups 1-8 A indicate the total number of valence electrons. • Main-group elements

30. 8 valence electrons 1 valence electron 3 valence electrons 2 valence electrons

31. The Periodic Table

32. Basic Facts • Dmitri Mendeleev, a Russian chemist, credited with invention in 1869 • Original purpose was to show trends • Changes were constantly being made to the table over time • Current standard periodic table has: • 118 elements • 18 Groups (columns) • 7 Periods (rows)

35. Areas of the Periodic Table (11.11) • Metals • Location • Trend • Properties • Non-metals • Location • Trend • Properties • Metalloids

36. Families (4.8) • Alkali Metals • Alkaline Earth Metals • Halogens • Noble Gases • Transition Metals

37. Alkali Metals • Highly reactive • Never found in elemental form in nature • Nearly all are silver colored • Soft metals • Low densities • Known for their vigorous reactions with water (video)

38. Alkaline Earth Metals • Silver in color • Soft metals • Have distinguishable flame color (except for Be and Mg) • Quite reactive with water, but not as much as alkali metals

39. Transition Metals • Name comes from their position on the periodic table • Magnetic under specific conditions

40. Halogens • Only group that contains all 3 states of matter • Produce a salt when it forms a compound with a metal • Only found in the environment as compounds or as ions because of high reactivity

41. Noble Gases • Little tendency to participate in chemical reactions because their outer shell of valence electrons is “full” • Because of low boiling point and melting points, have many practical uses.

43. Natural States of the Elements (4.9) • Noble Metals • Why are they Noble? • Examples: • Diatomic molecules • Definition: • What are the seven diatomic elements?

44. Trends related to Electron Configuration (11.11) • Electronegativity • Definition • Trend • Electron Affinity • Definition • Trend • Ionization Energy • Definition • Trend

45. Trends related to Electron Configuration (11.11) Atomic Radius Definition Trend Ionic Radius Definition Trend

48. Reactivity • Definition • Trend • Metallic Character • Non-Metallic Character • Practice Problems

50. Nuclear Chemistry