1. Ch. 1—Chemistry: An Introduction What is chemistry? Chemistry is the study of the ___________________ of substances and the changes they undergo. It began from “_______________”... the attempts of alchemists to change common metals into _________ through trial and error. Do you believe we can make gold? Why or why not? composition alchemy gold Half of chemistry in one sentence: “Atoms that don’t have enough electrons in the outer level will fight, barter, beg, make and break alliances, or do whatever they must to get the right number.” - Kean, Sam. The Disappearing Spoon. New York: Back Bay Publishing, 2010. Print.

2. How do we classify materials in chemistry? Elements cannot be ___________ down or _____________ into simpler substances by chemical means. Elements are the _________ forms of matter that can exists in normal laboratory conditions. Compounds are made up of ____ or ________ different elements ______________ bonded together. Compounds can only be broken down into simpler substances by ____________ ____________. Mixtures are a physical blend of two or more substances mixed together.” The parts can be separated by _____________ means or ____________ changes. physical brokenchanged simplest 2more chemically chemical reactions

3. Chemical Symbols Chemists use chemical symbols for the elements involved in a chemical reaction. The symbols are a shorthand way of representing the ______________. (See the Periodic Table for a list of all the symbols.) The first letter of the chemical symbol for an element is always _________________. The next letter, if needed, is _______________. Each capital letter in a formula, therefore, represents another element. Examples: ____, ____, Hg, ___, NaBr, ________, LiC 2 H 3 O 2 Some symbols come from _______ names: Au=Aurum (Gold) elements capitalized lowercase HNeS Latin H2OH2O

4. Chemical Reactions When writing chemical reactions, the substances that ___________ with each other are written on the _______ and are called “reactants”. The substances that are ____________ are written on the _______ and are called the “products.” Reactants  Products The “  ” symbol can be read as “_______” or “reacts to produce.” Example: 2H 2 + O 2  2H 2 O which means “____________________________________ ________________________________________________.” react left producedright yields two hydrogen moleculesplus one oxygen moleculeyieldstwo water molecules

5. Conservation of Mass During chemical (or physical) reactions, mass (or matter) is neither _____________ nor _________________. The mass of all the reactants _________ the mass of all the products. The ___________ of each kind of atom is the same. Sometimes it appears that the reactant and product masses are not equal, but a _______ was probably a reactant or product in the reaction, and that is making the difference! Example: 2H 2 + O 2  2H 2 O If 4 grams of hydrogen reacted with oxygen to produce 36 grams of water, how many grams of oxygen were used? _______ Notice that the ____ of H’s and O’s on each side is __________! createddestroyed equals number gas 32 #constant

6. Conservation of Mass CaCl 2 + Na 2 SO 4  CaSO 4 + 2NaCl mass before = mass after # atoms before = # atoms after

7. Atomic Theory and Structure The smallest particle of an ________________ is an atom. The atom is made up of three ________________ particles. The Theory of the Atom (1) ________________, a famous Greek teacher who lived in the 4th Century B.C., first suggested the idea of the atom. (2) ________ __________ came up with his solid sphere atomic theory based on the results of his experiments. (3)The proton has a ______ charge, and it was discovered in _________ by E. Goldstein. (4)The electron was discovered in _______ by J. J. Thomson by using a cathode ray tube. The electron has a _______ charge. It’s mass is much smaller than the other 2 subatomic particles, therefore it’s mass is usually ______________. Democritus John Dalton element subatomic 1897 (−)(−) ignored (+) 1886

8. Cathode Ray Tube

9. (4) Model: a ball of (+) charge containing a number of e - no ________________ often described as the “________ _______________” atom. Thomson nucleus plum pudding Atomic Models

10. The Nucleus (5) Discovered by Ernest ________________ in ________. He shot a beam of positively charged “alpha particles”, which are ___________ nuclei, at a thin sheet of ____ ____. 99.9% of the particles went right on through to the ______________. Some were slightly deflected. Some even ____________ ________ towards the source! This would be like shooting a cannon ball at a piece of tissue paper and having it bounce off. Rutherford1911 heliumgold foil detector bounced back

11. Rutherford’s Experiment

12. Most of the atom is more or less _________ ___________. The nucleus is very _________. (Stadium Analogy) The nucleus is very _________. (Large Mass ÷ Small Volume) The nucleus is ______________ charged. (5) Model: a ____________ of (+) charge surrounded by a number of e - no _____________ and no e - orbitals Rutherford nucleus neutrons Conclusions about the Nucleus empty space dense tiny positively

13. Nuclear Atomic Structure The atom is made up of 2 parts/sections: (1) The ______________ --- (in the center of the atom) (2) The ____________ _________ --- (surrounds the nucleus) nucleus electron cloud (p + & n 0 ) e − cloud

14. (6) The neutron does not have a charge. In other words, it is ________. It was discovered in ____ by James Chadwick. The neutron has about the same _________ as the proton. These three particles make up all the ____________________ in the Universe! There are other particles such as neutinos, positrons, and quarks, but are typically left for 2 nd year chemistry courses. neutral1932 mass visible matter Atomic Models

15. (7) Model: a nucleus of (+) charge that also contains ______________ nucleus is encircled by e - ’s located in definite orbits (or paths). e - ’s have ___________ energies in these orbits e - ’s do not lose energy as they orbit the nucleus Atomic Models Bohr neutrons fixed

16. Bohr Atomic Model

18. How to draw your own Bohr model The atomic number tells you how many electrons a neutral atom will contain. The first energy level can only fit _____electrons (like He). Each energy level beyond the first one can fit ______ electrons. (or 18 if it is after the middle block which is also called the d-block) 2, 8, 8, 18, 18… We will expand on this when we get to electron configurations. eight

19. Quantum Mechanical Model (8) Mechanical Model ( Wave Mechanical Model) no definite ____________ to the e - path (“fuzzy” cloud) orbits of e - ’s based on the _________________ of finding the e - in the particular orbital shape. Quantum shape probability

20. Quantum Mechanical Model

21. Schroedinger's Cat

23. Quantum Mechanical Model (present)

24. Counting Subatomic Particles in an Atom The atomic # of an element equals the number of ____________ in the nucleus. (This number is the whole number on the periodic table) The mass # of an element equals the sum of the _____________ and ______________ in the nucleus. (This number cannot be found on the periodic table) In a neutral atom, the # of protons = # of ______________. To calculate the # of neutrons in the nucleus, ______________ the ___________ # from the __________ #. protons neutrons electrons subtract atomicmass

25. Practice Problems (1)Find the # of e -, p + and n 0 for sodium. (mass # = 23) 2)Find the # of e -, p + and n 0 for uranium. (mass # = 238) 3) What is the atomic # and mass # for the following atom? # e - = 15; # n 0 = 16 Atomic # = 11 = # e - = # p + # neutrons = 23-11 = 12 Atomic # = 92 = # e - = # p + # neutrons = 238-92 = 146 Atomic # = 15 = # e - = # p + Mass # = p + + n 0 = 15+16 =31 The element is phosphorus!

26. Isotopes An isotope refers to atoms that have the same # of ___________, but they have a different # of ___________. Because of this, they have different _________ #’s (or simply, different ___________.) Isotopes are the same element, but the atoms weigh a different amount because of the # of ______________. Examples---> (1) Carbon-12 & Carbon-13 (2) Chlorine-35 & Chlorine-37 (The # shown after the name is the mass #.) For each example, the elements have identical ___________ #’s, (# of p + ) but different _________ #’s, (# of n 0 ). Another way to write the isotopes in shorthand is as follows: CCl 12 6 35 17 The top number is the ________ #, and the bottom # is the __________ number. Calculating the # n 0 can be found by _____________ the #’s! protons neutrons mass masses neutrons atomic mass atomic subtracting

27. Figure 3.10: Two isotopes of sodium.

28. More Practice Problems (1)Find the # e -, p + and n 0 for Xe-131. 2)Find the # e -, p + and n 0 for. 3) Write a shorthand way to represent the following isotope: # e - = 1 # n 0 = 0 # p + = 1 Cu 63 29 Atomic # = 54 = p + = e − n 0 = 131 − 54 = 77 Atomic # = 29 = p + = e − n 0 = 63 − 29 = 34 Atomic # = p + = e − = 1 mass # = n 0 + p + = 1+ 0 = 1 H-1 orH 1 1

29. Ions An atom can gain or lose electrons to become electrically charged. Cation = (___) charged atom created by ___________ e-’s. – Cations are ______________ than the original atom. – _____________ generally form cations. Anion = (___) charged atom created by _____________ e-’s. – Anions are ____________ than the original atom. – _______________ generally form anions. Practice Problems: Count the # of protons & electrons in each ion. a) Mg +2 p + = _____ e − = ______ b) F −1 p + = _____ e − = ______ +losing smaller Metals −gaining larger Nonmetals 1210 9

30. Atomic Mass Based on the relative mass of Carbon-12 which is exactly _______. 1 p + ≈ __ atomic mass unit (amu) 1 n 0 ≈ __ amu 1e - ≈ __ amu The atomic masses listed in the Periodic Table are a “weighted average” of all the isotopes of the element. 12 110 Weighted Average Practice Problems: (1) In chemistry semester grades are calculated using a weighted average of three category scores: Summative Assessments = 80% of your grade Formative Assessments = 10% of your grade Semester Exam = 10% of your grade If a student had the following scores, what would they receive for the semester? Summative Assessments = 3.5 Formative Assessments = 2.5 Semester Exam = 2.0

31. Weighted Average Step (1): Multiply each score by the % that it is weighted. Step (2): Add these products up, and that is the weighted average! 3.5 x.80 = 2.80 2.5 x.10 = 0.25 2.0 x.10 = 0.20 Add them up!! A “normal average” would be calculated by simply adding the raw scores together and dividing by 3… 3.5 + 2.5 + 2.0 = 8 ÷ 3 = 2.6 (C) + 3.25 (B+)