Chemistry I – Unit 2 Notes Te x tbook Correlation

1. Chemistry I – Unit 2 Notes Textbook Correlation Unit 2 consists of the following sections in the text: 4.1-4.3, 4.5-4.7, 4.10, 8.1, 11.1-11.9

2. 4.1 - The Elements • A. An element is: • B. Elements can exist as pure substances or as parts of compounds. • Examples: • 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. 4.3 – Dalton’s Atomic Theory John Dalton (1766-1844) – English Scientist page 78 Dalton’s Atomic Theory 1. Elements are made of tiny particles called atoms. 2. All atoms of a given element are identical (we now know that this is not exactly true) 3. The atoms of a given element are different from those of any other elements. 4. Atoms of one element can combine with atoms of other elements to form compounds. A given compound always has the same relative numbers and types of atoms. Examples: 5. Atoms are indivisible in chemical processes. That is, atoms are not created or destroyed in chemical reactions. A chemical reaction simply changes the way the atoms are grouped together.

4. 4.5 – The Structure of the Atom A. J. J. Thomson’s Plum Pudding Model (1890’s) described atoms as being:

5. B. Ernest Rutherford’s “Nuclear” Model (1911) of the atom described the structure of atoms as follows:

6. 1. Gold Foil Experiment (Rutherford)

7. 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

8. 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

9. 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:

10. 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:

11. E. A “Box” in the Periodic Table:

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

13. 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.

14. 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

15. Example 2

16. 11.1 Rutherford’s Atom A. His model leaves many questions about electrons unanswered 1. How are the electrons arranged? 2. How do they move? 3. Since the nucleus and the electrons are oppositely charged, why doesn’t the atom collapse?

17. 11.2 Electromagnetic Radiation – energy transmitted as a wave A. Parts of a wave 1. Wavelength is – 2. Frequency is –

18. B. Types of EM Radiation and wavelength.

19. 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:

20. 11.4 The Energy Levels of Hydrogen A. Electrons can only absorb quantized amounts of energy…this means ___________________________________________________________ B. With only one electron, a hydrogen atom is the simplest way to view what can happen when electrons get excited. C. How can Hydrogen produce photons with 4 different energies (colors) when it only has 1 electron to excite? (See Figure 11.13 on pg. 310)

21. 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 –

22. 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.

23. 11.8 Organization of Principle Energy Levels, Orbitals, and Sublevels. A. Principle energy levels contain sublevels, which in turn contain orbitals. B. Pauli exclusion principle:

24. 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. Hund’s Rule • 4. Valence Electrons • 5. Core Electrons • 6. Examples from pg. 319-320