Unit 1 review. Chapter 1, chapter , 2.4

Transcription

1 Unit 1 review Chapter 1, chapter , 2.4

2 The Organization of Matter Matter Mixtures: a) Homogeneous (Solutions) b) Heterogeneous Pure Substances Elements Compounds Atoms Nucleus Protons Quarks Electrons Neutrons Quarks

3 Pure Substances Each have a fixed composition Each a unique set of properties Mercury, Hg Elements a type of matter that cannot be broken down into two or more pure substances Chlorine gas, Cl 2 Silver metal, Ag Sodium metal, Na Compounds A pure substance that contains more that one element. Sodium Chloride, NaCl Water, H2O

4 Phase Differences Solid definite volume and shape; particles packed in fixed positions. Liquid definite volume but indefinite shape; particles close together but not in fixed positions Gas neither definite volume nor definite shape; particles are at great distances from one another Plasma high temperature, ionized phase of matter as found on the sun.

5 Properties of Substances Chemical properties Observed when substances take part in a chemical reaction A change that converts it to a new substance. Physical properties Observed without changing the chemical identity of a substance. Properties of Gold are: Melting point Boiling point Color Texture Density others Physical: Melting point of 1063 oc (intensive) Color gold (intensive) Amount in weight (extensive) Chemical: Gold can be stored in air without reacting chemically with oxygen

6 Separation of a Compound The Electrolysis of water Compounds must be separated by chemical means. With the application of electricity, water can be separated into its elements Reactant Water 2 H 2O Products Hydrogen + Oxygen 2 H2 + O2

7 Mixtures Contain two or more substances combined in such a way that each substance retains its chemical identity. Homogeneous Uniform mixture in which the composition is the same throughout. Granite Brass Heterogeneous Homogeneous Heterogeneous Copper Sulfate, heterogeneous Nonuniform, different in composition throughout.

8 Separation of a Mixture The constituents of the mixture retain their identity and may be separated by physical means.

9 Separation of a Mixture The components of dyes such as ink may be separated by paper chromatography.

10 Separation of a Mixture by Distillation

11 Measurements Scientific measurements are expressed in the metric system. You will need to review this system if you are unfamiliar with it. Pages 14 in you text.

12 Fundamental SI units (systeme international-units agreed upon by the science community) Mass is measured in grams Length is measured in meters Time is measure in seconds Temperature is measured in Kelvin Electric currents is measure in amperes Amount of substance is measured in moles All other units are derived from these basic units

13 Temperature Kelvin = celsius+ 273

14 Derived units Volume measured in cm3 Derived from finding the volume of a cube Length (cm) x width (cm) x height (cm) = cm3 Common unit (but not the SI) is liters The AP test will use liters or milliliters

15 Density the ratio of mass to volume. D = m/v SI unit: g/cm3 common in g/ml AP test uses: g ml-1

16 Significant Figures Every measurement carries with it a degree of uncertainty. This depends upon what instrument is being used to measure. We do not use the + in significant figures however it is understood that there is an uncertainty of at least one unit in the last digit place.

17 Rules of significant figures If it is a number 1-9 it is significant. If it is a zero between two numbers, it is significant. If it is a zero that tells how well something was measured, it is significant. (zeros at the end with a decimal in the number) If it is a zero that just tells how big or how small a number is, it is NOT significant. (zeros at the beginning and at the end with no decimal in the number)

18 Examples How many significant figures are in each? x 103 Answers

19 When multiplying or dividing significant figures: Count how many significant figures in each of the numbers being used. Then use the smallest amount of significant figures of all the numbers when reporting the answer. When adding or subtracting significant figures: Line up the decimals, the last place that is significant in both numbers is where you draw a line. Add or subtract the numbers, then look at the number just past the line if it is bigger or equal to 5 round up if not just drop after the line.

20 Examples Calculate and record the correct number of significant figures x x 10 23/1.0 x Answers x

21 Dalton s Atomic Theory (1808) John Dalton All matter is composed of extremely small particles called atoms Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties Atoms cannot be subdivided, created, or destroyed Atoms of different elements combine in simple whole-number ratios to form chemical compounds In chemical reactions, atoms are combined, separated, or rearranged

22 Modern Atomic Theory Several changes have been made to Dalton s theory. Dalton said: Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties Modern theory states: Atoms of an element have a characteristic average mass which is unique to that element.

23 Modern Atomic Theory #2 Dalton said: Atoms cannot be subdivided, created, or destroyed Modern theory states: Atoms cannot be subdivided, created, or destroyed in ordinary chemical reactions. However, these changes CAN occur in nuclear reactions

24 Discovery of the Electron In 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle. Cathode ray tubes pass electricity through a gas that is contained at a very low pressure.

25 Thomson s Atomic Model J.J. Thomson Thomson believed that the electrons were like plums embedded in a positively charged pudding, thus it was called the plum pudding model.

26 Mass of the Electron 1909 Robert Millikan determines the mass of the electron. Mass of the electron is x kg The oil drop apparatus

27 Conclusions from the Study of the Electron Cathode rays have identical properties regardless of the element used to produce them. All elements must contain identically charged electrons. Atoms are neutral, so there must be positive particles in the atom to balance the negative charge of the electrons Electrons have so little mass that atoms must contain other particles that account for most of the mass

28 Rutherford s Gold Foil Experiment Alpha particles are helium nuclei Particles were fired at a thin sheet of gold foil Particle hits on the detecting screen (film) are recorded

29 Try it Yourself! In the following pictures, there is a target hidden by a cloud. To figure out the shape of the target, we shot some beams into the cloud and recorded where the beams came out. Can you figure out the shape of the target?

30 The Answers Target #1 Target #2

31 Rutherford s Findings Most of the particles passed right through A few particles were deflected VERY FEW were greatly deflected Like howitzer shells bouncing off of tissue paper! Conclusion s: The nucleus is small The nucleus is dense The nucleus is positively charged

32 Atomic Particles

33 The Atomic Scale Helium-4 Most of the mass of the atom is in the nucleus (protons and neutrons) Electrons are found outside of the nucleus (the electron cloud) Most of the volume of the atom is empty space Image: User Yzmo Wikimedia Commons.

34 Atomic Number Atomic number (Z) of an element is the number of protons in the nucleus of each atom of that element. This identifies the atom.

35 Mass Number Mass number is the number of protons and neutrons in the nucleus of an isotope. Mass # = p+ + n0 1 8 Arseni 7 c Phosphor5 us Electrons mass are so much smaller than a proton and neutron that they don t Contribute much to the overall mass of the atom,therefore they are not Counted in the mass number of the atom.

36 Sulfur Silver Periodic Table Metals good conductors of heat and electricity Nonmetals- nonconductors Metalloids -semiconductors Periods-horizontal rows Families/groups-vertical columns

37 Periodic table Group names- elements in the same group will react similarly 1- alkali metals (except Hydrogen) -most reactive metals 2- alkaline earth metals transitions metals 17- halogens- most reactive nonmetals 18- noble gases (inert or unreactive) Main group elements- groups 1,2,13-18

38 Percent error Percentage error = accepted value experimental value accepted value

39 Unit 1.1 Chapters 2.3, 3.3 Crash Course: chapter 2

40 Isotopes Isotopes are atoms of the same element having different masses due to varying numbers of neutrons.

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42 Atomic Masses Atomic mass is the average of all the naturally isotopes of that element. Carbon =

43 Purpose of Mass Spectrometry Produces spectra of masses from the molecules in a sample of material, and fragments of the molecules. Used to determine the elemental composition of a sample the masses of particles and of molecules potential chemical structures of molecules by analyzing the fragments the identity of unknown compounds by determining mass and matching to known spectra the isotopic composition of elements in a molecule

44 Stages The ionizer converts some of the sample into ions. Mass analyzers separate the ions according to their mass-to-charge ratio. The detector records either the charge induced or the current produced when an ion passes by or hits a surface

45 Interpreting Mass Spectra The height of each peak is proportional to the amount of each isotope present (i.e. it s relative abundance). The m/z ratio for each peak is found from the accelerating voltage for each peak. Many ions have a +1 charge so that the m/z ratio is numerically equal to mass of the ion.

46 Calculating the relative atomic mass from mass spectrometry 1. Measure the height of each peak. 2. Calculate the percentage relative abundance % abundance = amount of isotope x 100 total amount of all isotopes 3. Calculate the average mass mass of isotope A x abundance + mass of isotope B x abundance 100

47 Atomic and molecular weights Average atomic mass is the weighted average of all isotopes of an atom H = amu amu=atomic mass unit O = amu Average atomic mass is on the periodic table and takes into account all isotopes of an atom and their abundance

48 Unit 1.2 Chapter 3.4 Crash course: chapter 2

49 The Mole (mol) 1 dozen = 12 1 gross = ream = mole = 6.02 x 1023 There are exactly 12 grams of carbon-12 in one mole of carbon-12.

50 Avogadro s Number 6.02 x 1023 is called Avogadro s Number in honor of the Italian chemist Amadeo Avogadro ( ). I didn t discover it. Its just named after me! Amadeo Avogadro

51 Molar mass 1 mol of book and 1 mol of feather Same number of items, NOT the same mass Molar mass: mass in 1 mole of a substance (g/mol or g mol-1) Molar mass = formula mass or average atomic mass 1 mol a carbon = 6.02 x atoms = g 1 mol NaOH = 6.02 x formulas = 40 g 1 mol CO = 6.02 x molecules = 28 g

52 Calculations with Moles: Converting grams to moles How many moles of lithium are in 18.2 grams of lithium? 18.2 g Li 1 mol Li 6.94 g Li =2.62 mol Li

53 Calculations with Moles: Converting moles to grams How many grams of lithium are in 3.50 moles of lithium? 3.50 mol Li 6.94 g Li 1 mol Li = 24.3 g Li

54 Calculations with Moles: Using Avogadro s Number How many atoms of lithium are in 3.50 moles of lithium? 3.50 mol 6.02 x 1023 atoms 1 mol = 2.11 x 1024 atoms

55 Calculations with Moles: Using Avogadro s Number How many atoms of lithium are in 18.2 g of lithium? 18.2 g Li 1 mol Li 6.94 g Li x 1023 atoms Li 1 mol Li (18.2)(6.022 x 1023)/6.94 =1.58 x 1024 atoms Li

56 Moles Moles will convert to grams using molar mass Moles will convert to atoms/molecules/compounds using Avogadro's number

57 Formula mass = mass in one chemical formula Add the mass of all atoms in the formula Calculating Formula Mass Calculate the formula mass of carbon dioxide, CO g + 2(16.00 g) = g

58 Formula mass = mass in one chemical formula Add the mass of all atoms in the formula Calculating Formula Mass Calculate the formula mass of carbon dioxide, CO g + 2(16.00 g) = g

59 Conversions How many grams are in 2.50 mol of oxygen gas? (2.50 mol) (32 g O2/ 1mole) 80.0 g

60 How many molecules are in 25.0 g of sulfuric acid? (25.0 g) (1 mol/98.1 g H2SO4) (6.02x1023 molecules/1 mol) 1.53 x molecules How many hydrogen atoms are in 25.0 g of sulfuric acid 2(1.53 x molecules)

61 Ibuprofen, C13H18O2 has a molar mass of g/mol. If a bottle of ibuprofen contains 33 g of it, how many moles of ibuprofen are in the bottle and how many molecules are there? 0.16 moles, 9.6 x 1022 molecules

62 Unit 1.3 Chapter 3.5 Crash course: chapter 2

63 Percent composition We can find the mass percent of each element in a compound. mass of element in compound x 100 total mass of compound

64 Calculating Percentage Composition Calculate the percentage composition of magnesium carbonate, MgCO3. Formula mass of magnesium carbonate: g g + 3(16.00 g) = g

65 Formulas Empirical formula: the lowest whole number ratio of atoms in a compound. Molecular formula: the true number of atoms of each element in the formula of a compound. molecular formula = (empirical formula)n [n = integer] molecular formula = C6H6 = (CH)6 empirical formula = CH

66 Formulas (continued) Formulas for ionic compounds (metals bonded to nonmetals) are ALWAYS empirical (lowest whole number ratio). Examples: MgCl2 Al2(SO4)3 K2CO3 NaCl Ionic compounds to not form molecules (they form crystals) so the formula doesn t show the exact number of atoms in the compound but instead a ratio of how they bond.

67 Formulas (continued) Formulas for molecular compounds (nonmetals bonded to nonmetals to form molecules) MIGHT be empirical (lowest whole number ratio). Molecular: H2O C6H12O6 C12H22O11 Empirical: H2O CH2O C12H22O11

68 Empirical Formula Determination 1. Base calculation on 100 grams of compound. 2. Determine moles of each element in 100 grams of compound Divide each value of moles by the smallest of the values. Multiply each number by an integer to obtain all whole numbers. 73.9% Hg and 26.1% Cl 73.9 g Hg and 26.1 g Cl 73.9/200.6 = mol Hg 26.1/35.5 = mol Cl 0.735/0.368 = /0.368 = 1 HgCl2

69 Make sure to go at least two places past decimal when finding molar mass If you get a 1.5:1 ratio, double both 3:2 If you get a 1.33 : 1 ratio, triple both 4:3 What about 1.25: 1?

70 Empirical Formula Determination Adipic acid contains 49.32% C, 43.84% O, and 6.85% H by mass. What is the empirical formula of adipic acid?

71 Empirical Formula Determination (part 2) Divide each value of moles by the smallest of the values. Carbon: Hydrogen: Oxygen:

72 Empirical Formula Determination (part 3) Multiply each number by an integer to obtain all whole numbers. Carbon: 1.50 x 2 3 Hydrogen: 2.50 x 2 5 Oxygen: 1.00 x 2 2 Empirical formula: C H O 3 5 2

73 Finding the Molecular Formula The empirical formula for adipic acid is C3H5O2. The molecular mass of adipic acid is 146 g/mol. What is the molecular formula of adipic acid? 1. Find the formula mass of C3H5O2 3(12.01 g) + 5(1.01) + 2(16.00) = g

74 Finding the Molecular Formula The empirical formula for adipic acid is C3H5O2. The molecular mass of adipic acid is 146 g/mol. What is the molecular formula of adipic acid? 2. Divide the molecular mass by the mass given by the emipirical formula. 3(12.01 g) + 5(1.01) + 2(16.00) = g

75 Finding the Molecular Formula The empirical formula for adipic acid is C3H5O2. The molecular mass of adipic acid is 146 g/mol. What is the molecular formula of adipic acid? 3. Multiply the empirical formula by this number to get the molecular formula. 3(12.01 g) + 5(1.01) + 2(16.00) = g (C3H5O2) x 2 = C6H10O4

76 example Empirical formula C3H4 Molecular mass = 121 amu What s the molecular formula? Mass of empirical formula = 40 amu 121/40 = 3.02 Molecular formula: C9H12