Intro to Honors Chemistry

1. Intro to Honors Chemistry A Review of the Summer Assignment…

2. Structure Determines Properties carbon dioxide carbon monoxide • composed of one carbon atom and two oxygen atoms • colorless, odorless gas • incombustible • does not bind to hemoglobin • composed of one carbon atom and one oxygen atom • colorless, odorless gas • burns with a blue flame • binds to hemoglobin • The properties of matter are determined by the atoms and molecules that compose it

3. Scientific Method Procedure designed to test an idea Tentative explanation of a single or small number of observations General explanation of natural phenomena Careful noting and recording of natural phenomena Generally observed occurence in nature

4. Classifying Matterby Physical State • Matter can be classified as solid, liquid, or gas based on the characteristics it exhibits

6. Changes in Matter • Changes that alter the state or appearance of the matter without altering the composition are called physical changes • Changes that alter the composition of the matter are called chemical changes • during the chemical change, the atoms that are present rearrange into new molecules, but all of the original atoms are still present

7. Properties of Matter • Physical properties are the characteristics of matter that can be changed without changing its composition • characteristics that are directly observable • Chemical properties are the characteristics that determine how the composition of matter changes as a result of contact with other matter or the influence of energy • characteristics that describe the behavior of matter

8. Energy Changes in Matter • Changes in matter, both physical and chemical, result in the matter either gaining or releasing energy • Energy is the capacity to do work • Work is the action of a force applied across a distance • a force is a push or a pull on an object • electrostatic force is the push or pull on objects that have an electrical charge

9. Conversion of Energy • You can interconvert kinetic energy and potential energy • Whatever process you do that converts energy from one type or form to another, the total amount of energy remains the same • Law of Conservation of Energy

10. Spontaneous Processes • Materials that possess high potential energy are less stable • Processes in nature tend to occur on their own when the result is material with lower total potential energy • processes that result in materials with higher total potential energy can occur, but generally will not happen without input of energy from an outside source

11. Counting Significant Figures • All non-zero digits are significant • 1.5 has 2 sig. figs. • Interior zeros are significant • 1.05 has 3 sig. figs. • Leading zeros are NOT significant • 0.001050 has 4 sig. figs. • 1.050 x 10−3

12. Counting Significant Figures • Trailing zeros may or may not be significant • Trailing zeros after a decimal point are significant • 1.050 has 4 sig. figs. • Trailing zeros before a decimal point are significant if the decimal point is written • 150.0 has 4 sig. figs. • Zeros at the end of a number without a written decimal point are ambiguous and should be avoided by using scientific notation • if 150 has 2 sig. figs. then 1.5 x 102 • but if 150 has 3 sig. figs. then 1.50 x 102

13. Example 1.5: Determining the Number of Significant Figures in a Number How many significant figures are in each of the following? 0.04450 m 5.0003 km 10 dm = 1 m 1.000 × 105 s 0.00002 mm 10,000 m 4 sig. figs.; the digits 4 and 5, and the trailing 0 5 sig. figs.; the digits 5 and 3, and the interior 0’s infinite number of sig. figs., exact numbers 4 sig. figs.; the digit 1, and the trailing 0’s 1 sig. figs.; the digit 2, not the leading 0’s Ambiguous, generally assume 1 sig. fig.

14. Multiplication and Division with Significant Figures • When multiplying or dividing measurements with significant figures, the result has the same number of significant figures as the measurement with the lowest number of significant figures 5.02 × 89.665 × 0.10 = 45.0118 = 45 3 sig. figs. 5 sig. figs. 2 sig. figs. 2 sig. figs. 5.892 ÷ 6.10 = 0.96590 = 0.966 4 sig. figs. 3 sig. figs. 3 sig. figs.

15. Addition and Subtraction with Significant Figures • When adding or subtracting measurements with significant figures, the result has the same number of decimal places as the measurement with the lowest number of decimal places

16. Both Multiplication/Division and Addition/Subtraction with Significant Figures • When doing different kinds of operations with measurements with significant figures, do whatever is in parentheses first, evaluate the significant figures in the intermediate answer, then do the remaining steps 3.489 × (5.67 – 2.3) = 2 dp 1 dp 3.489 × 3.37 = 12 4 sf 1 dp & 2 sf 2 sf

17. Accuracy vs. Precision Suppose three students are asked to determine the mass of an object whose known mass is 10.00 g The results they report are as follows Looking at the graph of the results shows that Student A is neither accurate nor precise, Student B is inaccurate, but is precise, and Student C is both accurate and precise.

18. Problem Solving and Dimensional Analysis Arrange conversion factors so the starting unit cancels arrange conversion factors so the starting unit is on the bottom of the first conversion factor May string conversion factors so you do not need to know every relationship, as long as you can find something else the starting and desired units are related to

19. Order of Magnitude Estimations • Using scientific notation • Focus on the exponent on 10 • If the decimal part of the number is less than 5, just drop it • If the decimal part of the number is greater than 5, increase the exponent on 10 by 1 • Multiply by adding exponents, divide by subtracting exponents

20. Law of Conservation of Mass • In a chemical reaction, matter is neither created nor destroyed • Total mass of the materials you have before the reaction must equal the total mass of the materials you have at the end • total mass of reactants = total mass of products Antoine Lavoisier 1743-1794

21. Law of Definite Proportions Joseph Proust 1754-1826 • All samples of a given compound, regardless of their source or how they were prepared, have the same proportions of their constituent elements

22. Law of Multiple Proportions John Dalton 1766-1844 • When two elements (call them A and B) form two different compounds, the masses of B that combine with 1 g of A can be expressed as a ratio of small, whole numbers

23. Some Notes on Charge Two kinds of charge called + and – Opposite charges attract + attracted to – Like charges repel + repels + – repels – To be neutral, something must have no charge or equal amounts of opposite charges

24. Thomson’s Experiment Investigate the effect of placing an electric field around tube 1. charged matter is attracted to an electric field 2. light’s path is not deflected by an electric field +++++++++++ Anode Cathode (+) (-) ------------- - + Power Supply

25. Thomson’s Conclusions, cont’d • Thomson believed that these particles were therefore the ultimate building blocks of matter • “We have in the cathode rays matter in a new state, a state in which the subdivision of matter is carried very much further . . . a state in which all matter . . . is of one and the same kind; this matter being the substance from which all the chemical elements are built up.” • These cathode ray particles became known as electrons

26. Electrons Electrons are tiny, negatively charged particles found in all atoms Cathode rays are made of streams of electrons The electron has a charge of −1.60 x 1019 C The electron has a mass of 9.1 x 10−28 g

27. Radioactivity In the late 1800s, Henri Becquerel and Marie Curie discovered that certain elements would constantly emit small, energetic particles and rays These energetic particles could penetrate matter Ernest Rutherford discovered that there were three different kinds of emissions alpha, a, rays made of particles with a mass 4x H atom and + charge beta, b, rays made of particles with a mass ~1/2000th H atom and – charge gamma, g, rays that are energy rays, not particles Marie Curie 1867-1934

28. Rutherford’s Experiment • How can you prove something is empty space? • Put something through it! • use large target atoms • use very thin sheets of target so it will not absorb “bullet” • use very small particle as bullet with very high energy • but not so small that electrons will affect it • Bullet = alpha particles, target atoms = gold foil • a particles have a mass of 4 amu & charge of +2 c.u. • gold has a mass of 197 amu & is very malleable

30. Rutherford’s Results • Over 98% of the a particles went straight through • About 2% of the a particles went through but were deflected by large angles • About 0.005% of the a particles bounced off the gold foil • “...as if you fired a 15” cannon shell at a piece of tissue paper and it came back and hit you.”

31. Rutherford’s Conclusions • Atom mostly empty space • because almost all the particles went straight through • Atom contains a dense particle that is small in volume compared to the atom but large in mass • because of the few particles that bounced back • This dense particle is positively charged • because of the large deflections of some of the particles

32. Plum Pudding Atom • • • • If atom was like a plum pudding, all the a particles should go straight through • • • • • • • • • • • • • • • • • • A few of the a particles do not go through Nuclear Atom . Almost all a particles go straight through . Some a particles go through, but are deflected due to +:+ repulsion from the nucleus .

33. Rutherford’s Interpretation –the Nuclear Model The atom contains a tiny dense center called the nucleus the amount of space taken by the nucleus is only about 1/10 trillionth the volume of the atom The nucleus has essentially the entire mass of the atom the electrons weigh so little they give practically no mass to the atom The nucleus is positively charged the amount of positive charge balances the negative charge of the electrons The electrons are dispersed in the empty space of the atom surrounding the nucleus

34. Structure of the Nucleus • Rutherford proposed that the nucleus had a particle that had the same amount of charge as an electron but opposite sign – these particles are called protons • based on measurements of the nuclear charge of the elements • protons are subatomic particles found in the nucleus with a charge = +1.60 x 1019 C and a mass = 1.67262 x 10−24 g • Because protons and electrons have the same amount of charge, for the atom to be neutral there must be equal numbers of protons and electrons

35. Some Problems • How could beryllium have four protons stuck together in the nucleus? • shouldn’t they repel each other? • If a beryllium atom has four protons, then it should weigh 4 amu; but it actually weighs 9.01 amu! Where is the extra mass coming from? • each proton weighs 1 amu • remember, the electron’s mass is only about 0.00055 amu and Be has only four electrons – it can’t account for the extra 5 amu of mass

36. There Must Be Something Else! • To answer these questions, Rutherford and Chadwick proposed that there was another particle in the nucleus – it is called a neutron • Neutrons are subatomic particles with a mass = 1.67493 x 10−24 g and no charge, and are found in the nucleus • 1 amu • slightly heavier than a proton • no charge

38. Elements • Each element has a unique number of protons in its nucleus • The number of protons in the nucleus of an atom is called the atomic number • the elements are arranged on the Periodic Table in order of their atomic numbers • Each element has a unique name and symbol • symbol either one or two letters • one capital letter or one capital letter and one lowercase letter

40. The Periodic Table of the Elements The atomic number tells you how many protons are in the nucleus and how many electrons are in the atom Some symbols are one capital letter, like C, S, and I. Others are two letters, and the second is lowercase, like Br and Sr Some symbols come from the element ‘s name, like C for carbon. Others come from the Latin name of the element, like Au for gold (aurum) and Cu for copper (cuprium)

41. Structure of the Nucleus Soddy discovered that the same element could have atoms with different masses, which he called isotopes there are two isotopes of chlorine found in nature, one that has a mass of about 35 amu and another that weighs about 37 amu The observed mass is a weighted average of the weights of all the naturally occurring atoms the percentage of an element that is one isotope is called the isotope’s natural abundance the atomic mass of chlorine is 35.45 amu

42. Isotopes All isotopes of an element are chemically identical undergo the exact same chemical reactions All isotopes of an element have the same number of protons Isotopes of an element have different masses Isotopes of an element have different numbers of neutrons Isotopes are identified by their mass numbers, which is the sum of all the protons and neutrons in the nucleus

43. Isotopes • Atomic number • Number of protons • Z • Mass Number • Protons + neutrons • whole number • A • Abundance = relative amount found in a sample

44. Neon Symbol Number of Protons Number of Neutrons A, Mass Number Percent Natural Abundance Ne-20 or 10 10 20 90.48% Ne-21 or 10 11 21 0.27% Ne-22 or 10 12 22 9.25%

45. Example 2.3b: How many protons, electrons, and neutrons are in an atom of ? 52 Cr 24 52 Cr 24 atomic number symbol # p+ # e− atomic & mass numbers symbol # n0 Given: Find: therefore A = 52, Z = 24 # p+, # e−, # n0 Conceptual Plan: Relationships: in neutral atom, # p+ = # e- mass number = # p+ + # n0 Solution: Z = 24 = # p+ # e− = # p+ = 24 A = Z + # n0 52 = 24 + # n0 28 = # n0 Check: for most stable isotopes, n0≥ p+

46. Practice – Complete the table

47. Practice – Complete the table

48. Reacting Atoms When elements undergo chemical reactions, the reacting elements do not turn into other elements Statement 4 of Dalton’s Atomic Theory This requires that all the atoms present when you start the reaction will still be there after the reaction Because the number of protons determines the kind of element, the number of protons in the atom does not change in a chemical reaction However, many reactions involve transferring electrons from one atom to another

49. Charged Atoms When atoms gain or lose electrons, they acquire a charge Charged atoms or groups of atoms are called ions When atoms gain electrons, they become negatively charged ions, called anions When atoms lose electrons, they become positively charged ions, called cations

50. Ions and Compounds • Ions behave much differently than the neutral atoms • e.g., the metal sodium, made of neutral Na atoms, is highly reactive and quite unstable; however, the sodium cations, Na+, found in table salt are very nonreactive and stable • Because materials such as table salt are neutral, there must be equal amounts of charge from cations and anions in them