Ch. 1—Chemistry: An Introduction

1. composition alchemy Ch. 1—Chemistry: An Introduction gold 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? 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. broken changed simplest 2 more chemically chemical reactions physical physical

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, ________, LiC2H3O2 • Some symbols come from _______ names: Au=Aurum (Gold) elements capitalized lowercase H Ne S H2O Latin

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 “reactstoproduce.” • Example: 2H2 + O2 2H2O • which means “____________________________________ • ________________________________________________.” react left produced right yields two hydrogen molecules plus one oxygen molecule yields two 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: 2H2 + O2 2H2O • 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 __________! created destroyed equals number gas 32 # constant

6. Conservation of Mass CaCl2 + Na2SO4 CaSO4 + 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 ______________. element subatomic Democritus John Dalton (+) 1886 1897 (−) ignored

8. Cathode Ray Tube

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

10. Rutherford 1911 helium gold foil 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. detector bounced back

11. Rutherford’s Experiment

12. Conclusions about the Nucleus empty space tiny • 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 dense positively Rutherford nucleus neutrons

13. nucleus 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) electron cloud (p+ & n0) e− cloud

14. AtomicModels neutral 1932 mass (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 2nd year chemistry courses. visible matter

15. AtomicModels Bohr • (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 neutrons fixed

16. Bohr Atomic Model

17. 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 protons neutrons electrons subtract atomic mass

25. Practice Problems Atomic # = 11 = # e- = # p+ # neutrons = 23-11 = 12 • Find the # of e-, p+ and n0 for sodium. (mass # = 23) • Find the # of e-, p+ and n0 for uranium. (mass # = 238) 3) What is the atomic # and mass # for the following atom? # e- = 15; # n0 = 16 Atomic # = 92 = # e- = # p+ # neutrons = 238-92 = 146 Atomic # = 15 = # e- = # p+ Mass # = p+ + n0 = 15+16 =31 The element is phosphorus!

26. Isotopes protons neutrons mass • 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 n0). • Another way to write the isotopes in shorthand is as follows: masses neutrons atomic mass 12 C 35 Cl 6 17 mass atomic The top number is the ________ #, and the bottom # is the __________ number. Calculating the # n0 can be found by _____________ the #’s! subtracting

27. Figure 3.10: Two isotopes of sodium.

28. More Practice Problems Atomic # = 54 = p+ = e− n0 = 131 − 54 = 77 • Find the # e-, p+ and n0 for Xe-131. • Find the # e-, p+ and n0 for . 3) Write a shorthand way to represent the following isotope: # e- = 1 # n0 = 0 # p+ = 1 63 Cu 29 Atomic # = 29 = p+ = e− n0 = 63 − 29 = 34 Atomic # = p+ = e− = 1 mass # = n0 + p+ = 1+ 0 = 1 H-1 or 1 H 1

29. Ions + losing • 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− = ______ smaller Metals − gaining larger Nonmetals 12 10 9 10

30. Atomic Mass 12 1 1 0 • Based on the relative mass of Carbon-12 which is exactly _______. • 1 p+ ≈ __ atomic mass unit (amu) 1 n0 ≈ __ amu 1e- ≈ __ amu • The atomic masses listed in the Periodic Table are a “weighted average” of all the isotopes of the element. 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+)