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AP Chemistry Topic#1

Foundation Notes

Objectives

●

Table of Contents

1. Chemistry: An Overview

2. Scientific Method

3. Units of Measurement

4. Uncertainty in Measurement

5. Significant Figures and Calculations

6. Dimensional Analysis

7. Temperature

8. Density

9. Classification of Matter

10. The Early History of Chemistry

11. Fundamental Chemical Laws

12. Early Experiments to Characterize the Atom

13. Modern View of Atomic Structure: An Introduction

Chemistry: An Overview Topic#1

● Observations

o Macroscopic

□ Matter we can see for ourselves

o Microscopic

□ Matter we cannot see without assistance

● Atom, molecules, ions, etc.

o Quantitative

□ Measurement

● Number and unit

o Qualitative

□ Using your senses

● Chemicals

o Chemical elements

□ All of the discovered elements in the periodic table

□ Diatomic elements

● When pure, exist as two bonded atoms

● H2, N2, O2, F2, Cl2, Br2, and I2

o Compounds

□ Two or more elements chemically bonded together

o All elements and compounds are pure substances

□ Contains only the element or compound, no impurities

● Science

o A process for understanding nature

□ Make observations (data)

□ Predict (hypothesis)

□ Test prediction (experiment)

Scientific Method Topic#2

● Steps

o Make observation

o Hypothesis

o Experiment

o Theory (model)

□ Set of observations agreeing with various observations

□ Overall explanation of phenomena

● Observation

o Witnessed and can be recorded

● Theory

o An interpretation

▪ A possible explanation

▪ Subject to change with more information

o Tested by using to make a prediction

▪ Perform experiment to see whether results match

o Natural Law

□ Observations apply to many different systems

□ Generally observed behavior

● Law of conservation of Mass

o The total mass of materials is not affected by a chemical change

□ Difference between natural law and theory

● A law summarizes what occurs, while a theory (model) attempts to explain why.

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Units of Measurements Topic#3

● SI Units

|The Fundamental SI Units |

| Physical Quantity |Name of Unit |Abbr |

|Mass |kilogram |kg |

|Length |meter |m |

|Time |second |s |

|Temperature |Kelvin |K |

|Amount of substance |mole |mol |

o Based on the metric system

□ Derived units

● Made from SI units

o Volume

▪ 1 L = (1dm)3 = (10cm)3 = 1000cm3

• 1ml = 1cm3

• 1L = 1000mL = 1000cm3

▪ Measured

• Graduated cylinder

• Pipet

• Buret

• Volumetric flask

o Area

▪ 1 m2

● Prefixes

|Metric Prefixes |

|Prefix |Symbol |Meaning |Exponential |

| | | |Notation |

|exa |E |1,000,000,000,000,000,000 |1018 |

|peta |P |1,000,000,000,000,000 |1015 |

|tera |T |1,000,000,000,000 |1012 |

|giga |G |1,000,000,000, |109 |

|mega |M |1,000,000 |106 |

|kilo |k |1,000 |103 |

|hecto |h |100 |102 |

|deca |da |10 |101 |

|- |- |1 |100 |

|deci |d |0.1 |10-1 |

|centi |c |0.01 |10-2 |

|milli |m |0.001 |10-3 |

|micro |μ |0.000 001 |10-6 |

|nano |n |0.000 000 001 |10-9 |

|pico |p |0.000 000 000 001 |10-12 |

|femto |f |0.000 000 000 000 001 |10-15 |

|atto |a |0.000 000 000 000 000 001 |10-18 |

*most commonly used prefixes in chemistry

o Mass

□ The amt of matter in an object

● Does not change

o Weight

□ Gravities effect on a mass

● Changes based on gravitational pull

Uncertainty in Measurement Topic#4

● Uncertain digit

o The last digit measured

□ Always estimated

● Even in electronic devices

□ Take

● Person Results of Measurement

1. 20.15 mL

2. 20.14 mL

3. 20.16 mL

4. 20.17 mL

5. 20.16 mL

The first three numbers (20.1) are the same, but last digit varies

Called uncertain digit

□ Measurements

● Include all certain digits and one uncertain digit

● All have uncertainty

● Significant digits

o Sample Exercise 1.1 – Uncertainty in Measurement

In analyzing a sample of polluted water, chemist measured out a 25.00-mL water sample with a pipet. At another point in the analysis, the chemist used a graduated cylinder to measure 25 mL of a soln. What is the difference between the measurements 25.00 mL and 25 mL?

● Precision vs. Accuracy

o Accuracy

□ Agreement between a particular value and the true value

o Precision

□ Degree of agreement among several measurements of the same quantity

□ Errors

● Random

o Likely to occur with the same rate in any extreme

● Systematic

o Occurs in the same place all the time

o Sample Exercise 1.2 – Precision and Accuracy

To check the accuracy of a graduated cylinder, a student filled the cylinder to the 25-ml mark using water delivered from a buret and then read the volume delivered.

Trial Volume – Grad Cylinder Volume – Buret

1 25-mL 26.54-mL

2 25-mL 26.51-mL

3 25-mL 26.60-mL

4 25-mL 26.49-mL

5 25-mL 26.57-mL

Average 25-mL 26.54-mL

Is the graduated cylinder accurate? Is the graduated cylinder precise?

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Significant Figures and Calculations Topic#5

● Sig Fig Rules

1. Nonzero integers. Always count as sig figs.

2. Zeros. Count in between nonzero and at the end of a number with a DECIMAL point. Leading zeros never count.

3. Exact Numbers. Not used in calculating the number of sig figs in a data point. Include numbers from counting (10 experiments), defined in a formula (2πr2, the 2), from definitions (2.54cm = 1 inch).

o Sample Exercise 1.3 – Sig Figs

Give the number of sig figs for the following results.

(a) A student’s extraction procedure on tea yields 0.0105g of caffeine.

(b) A chemist records a mass of 0.050080g in an analysis.

(c) In an experiment a span of time is determined to be 8.050x10-3s.

● Sig Fig Rules for calculated numbers

1. Multiplication and division

□ The result has the same number of sig figs as the least precise data point.

● 4.56 x 1.4 = 6.38, the 1.4 has 2 sig figs so answer is 6.4

2. Addition and subtraction

□ The result has the same number of decimal places as the least precise data point.

● 12.11 + 18.0 + 1.013 = 31.123 (18.0 has 1 decimal place) = 31.1

● Rules for rounding

1. In a calculation, round at the end not during. Carry an extra digit or two.

2. Removing a digit

(a) Less than 5, then preceding digit stays the same, 1.33 becomes 1.3

(b) Equal or greater than five, preceding digit is increased by 1. 1.36 becomes 1.4.

3. Use only the first number to the right of the last significant figure

o Sample Exercise 1.4 – Significant Figures in Mathematical Operations

Carry out the following mathematical operations, and give each result with the correct number of sig figs.

(a) 1.05x10-3 ÷ 6.135

(b) 21 – 13.8

(c) As part of a lab assignment to determine the value of the gas constant (R), a student measured the pressure (P), volume (V), and temperature (T) for a sample of gas, where R = PV/T. The following values were obtained: P = 2.560, T = 275.15, and V = 8.8. Calculate R to the correct number of sig figs.

Visual Questions

What is the length of the pencil in the figure if the scale reads in centimeters? How many significant figures are there in this measurement?

An oven thermometer with a circular scale reading degrees Fahrenheit is shown. What temperature does the scale indicate? How many significant figures are in the measurement?

a) How many significant figures should be reported for the volume of the metal bar shown below? (b) If the mass of the bar is 104.7 g, how many significant figures should be reported when its density is calculated using the calculated volume?

Dimensional Analysis Topic#6

● Unit factor method/dimensional analysis

o Converting from one system of units to another

□ Conversion factor is made from an equivalence statement

● 1 in = 2.54 cm

o 1 in/2.54 cm or 2.54 cm/ 1 in (both equal to 1)

□ English to metrics equivalents

● Length

o 1 m = 1.094 yd

o 2.54 cm = 1 in

● Mass

o 1 kg = 2.205 lbs

o 453.6 g = 1 lb

● Volume

o 1 L = 1.06 qt

o 1 ft3 = 28.32 L

o Sample Exercise 1.5 – Unit Conversions I

A pencil is 7.00in long. What is its length in centimeters? Ans: 17.8cm

|Converting from One Unit to Another |

|To convert from one unit to another, use the equivalence statement that relates the two units. |

|Derive the appropriate unit factor by looking at the direction of the required change (to cancel the unwanted units). |

|Multiply the quantity to be converted by the unit factor to give the quantity with the desired units |

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o Sample Exercise 1.6 – Unit Conversion II

You want to order a bicycle with a 25.5in frame, but the sizes in the catalog are given only in centimeters. What size should you order? Ans: 64.8cm

o Sample Exercise 1.7 – Unit Conversion III

A student entered a 10.0km run. How long is the run in miles? Ans: 6.22 mi

o Sample Exercise 1.8 – Unit Conversion IV

The speed limit on many highways in the U.S. is 55mi/hr. What number would be posted in km/hr? Ans: 88 km/hr

o Sample Exercise 1.9 – Unit Conversion III

A Japanese car is advertised as having a gas mileage of 15lm/L. Convert this rating to mi/gal. Ans: 35 mi/gal

Temperature Topic#7

● The average KE of the particles in a sample

o KE of the particles is related to the speed of the particles in the sample

□ Some are slower/faster than average

● Bell curve

● Systems

o Celsius

□ bp - 100°C

□ fp - 0°C

o Kelvin

□ bp – 373K

□ fp – 273K

o Fahrenheit

□ bp - 212°F

□ fp - 32°F

● Converting between systems

o TK = TC + 273.15

o TC = TK -273.15

o TC = (TF -32)

1.8

o TF = 1.8TC + 32

o Sample Exercise 1.10 – Temperature I

Normal body temperature is 98.6°F. Convert this T to °C and K. Ans: 37°C and 310.2K

o Sample Exercise 1.11 – Temperature II

One interesting feature of the Celsius and Fahrenheit scales is that -40°C and -40°F represent the same T. Verify this true.

o Sample Exercise 1.12 – Temperature III

Liquid nitrogen, which is often used as a coolant for low-T experiments, has a bp of 77K. What is this T on the Fahrenheit scale? Ans: -321°F

Density Topic#8

● The ratio between the mass and volume of an object.

o By definition, matter is anything with a mass and a volume

□ Density is a property of matter

o Density = mass/volume = m/V

□ Solid

● g/cm3

□ Liquid

● g/mL (or g/cm3)

□ Gas

● g/L

o Sample Exercise 1.13 – Determining Density

A chemist, trying to identify the main component of a compact disc cleaning fluid, finds that 25.00cm3 of the substance has a mass of 19.625g at 20°C. The following are the names and densities of the compounds that might be the main component.

Compound Density (g/cm3)

Chloroform 1.492

Diethyl ether 0.714

Ethanol 0.789

Isopropyl alcohol 0.785

Toluene 0.867

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Classification of Matter Topic#9

● Matter

o Anything with a mass and volume (occupying space)

● States/Phases

o Solid

□ Rigid; fixed V and shape

□ Not compressible

o Liquid

□ Assumes shape of container; fixed V but not shape

□ Not compressible

o Gas

□ Assumes shape and V of container; no fixed V or shape

□ Highly compressible

● Matter Changes

o Physical (Δ)change

□ A change of phase

o Chemical (Δ) change

□ A change of one substance into another with different chemical properties

Does the diagram represent a chemical or physical change? How do you know?

● Types of Matter

o Pure

□ Constant composition

□ Cannot be separated by physical means

□ Elements

● Around 104 listed on Periodic Table

● Each specific element has a the same number of protons and electrons

o Isotopes of an atom have same number of protons and electrons but vary in number of neutrons

□ Compounds

● Has a definable formula (constant composition)

● Two or more chemically combined different atoms

● Chemical change

o Given substances turn into new substances with new, different substances

o Only method to separate elements in compounds

▪ Electrolysis

• H2O

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o Mixtures

□ Homogeneous (solutions)

● Uniform composition (visibly indistinguishable parts)

● Solute

o Part dissolved

o Less proportion

● Solvent

o Part dissolving

o Greater proportion

● Concentration

o Ratio of solute to solvent

□ Heterogeneous

● Non-uniform composition (visibly distinguished parts)

● Variable composition

□ Separation techniques of mixtures

● Physical change

o Change in state/phase of a substance

▪ s → l(melting), l → s (freezing), l → g (vaporization), g → l (condensation), s → g (sublimation, and g → s (deposition)

o Boiling a solution (boiling is vaporization at bp)

▪ solvent: l → g

• leaves solute behind

o Distillation

▪ Boiling a solution of liquids

▪ Depends on differences in volatilities

• Volatile

o Ease of vaporization

▪ Liquid with lowest bp will vaporize first

• Then next lowest

• Crude oil

o Filtration

▪ Separates a solid from a liquid

• Filter paper

▪ Separates small solids from large solids

• Screening

o Chromatography

▪ Two states of matter

• Stationary (solid) and mobile (liquid or gas)

• Separation occurs because the components of the mixture have different affinities for the two phases

• Component has a high affinity for mobile phase will move quickly through the chromatographic system

o The other component moves slower

▪ Paper chromatography

• Porous paper

o Stationary phase

• Solvent

o Mobile phase

▪ Travels up paper like a wick

▪ Components of mixture move at different speeds

• Drop of mixture is placed on stationary phase, then stationary phases is dipped in mobile phase

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Which of the figures represents (a) a pure element, (b) a mixture of two elements, (c) a pure compound, (d) a mixture of an element and a compound? (More than one picture might fit each description.)

The Early History of Chemistry Section#10

• Greeks first to try to explain why chemical changes occur

o Fire, earth, water, and air

o Demokritos – atomos, ultimate particles

o Next 2000 years dominated by pseudoscience – alchemy

o 16th century, foundations of modern science laid by the development of systematic metallurgy by a German, Georg Bauer, and medicinal use of minerals by Paracelsus

• Robert Boyle

o First “chemist” to perform quantitative experiments

o Relationship between pressure and volume of air

o “The Skeptical Chymist” (1661)

□ Born quantitative chemistry & physics

□ Experimental definition of an element (an element unless it could be broken down into two or more simpler substances).

• Joseph Priestley (1733-1804)

o Discovered oxygen

Fundamental Chemical Laws Section#11

• Antoine Lavoisier

o Law of conservation of mass

o Elementary Treatise on Chemistry

□ First chemistry textbook

□ Executed during French revolution by guillotine

• Joseph Proust

o Law of definite proportions

□ A given compound always contains exactly the same proportions of elements by mass

• John Dalton

o Law of multiple proportions

□ When two elements form a series of compounds, the ratios of the masses of the second element that combine with 1gram of the first element can always be reduced to small whole numbers.

o Sample Exercise 2.1 – Illustrating the Law of Multiple Proportions

The following data were collected for several compounds of nitrogen and oxygen.

Mass of N combining with 1g of O

Compound A 1.759g

Compound B 0.8750

Compound C 0.4375

Show how these data illustrate the law of multiple proportions

o Published A New System of Chemical Philosophy

□ Each element – tiny particles, atoms

□ Each element’s atoms, identical

□ Other element’s atoms different

□ Chemical compounds – two different atoms combine. Compound always has the same number and ratio of elements.

□ Chemical reaction is a rearrangement of atoms (changes in bonding). Atoms are not changed themselves.

□ 1st table of atomic masses (atomic weights)

• Gave H=1 and O=8, but was wrong

• Amadeo Avagadro along with Gay-Lussac

o Avagadro’s hypothesis

□ At same temperature, equal volumes of gases contain same number of particles

• 2H2+O2 → 2H2O

• 2 volumes of H2 react with 1 volume of O2 to produce 2 volumes of H2O

• 2 molecules of H2 react with 1 molecule of O2 to produce 2 molecules of H2O

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In a series of experiments, a chemist prepared three different compounds that contain only iodine and fluorine and determined the mass of each element in each compound. a) Calculate the mass of fluorine per gram of iodine in each compound. (b) How do the numbers in part (a) support the atomic theory?

Early Experiments to Characterize the Atom Section#12

• Electron

o J.J. Thomson (1900)

□ Cathode ray tube

□ Electron

• Charge to mass ratio

o e/m = -1.76x108 C/g

□ All atoms contain electrons (negative charges)

• Thus some positive charge

□ Plum pudding model

• Diffuse cloud of positive “charge” with the negative electrons embedded in it

o Robert Millikan (University of Chicago)

□ Oil drop experiment

• Charged electrons using gamma (x-rays) radiation to ionized

o Determined magnitude of charge

o Found mass of an e- to be 9.11x10-31 kg

• Radioactivity

o Henri Becquerel

□ Discovered radioactivity

• Uranium mineral left photographic image in the absence of light

o Ernest Rutherford

□ Discovered nucleus & proton

□ Gold foil experiment – shot α ( 42He2+)particles at gold foil

• To test Thomson’s plum pudding model

o Expected alpha particles to travel through foil with only minor deflections

o Most of the α particles went right through gold foil, but some were deflected at large angles

o Concluded Thomson’s model was incorrect

□ Nuclear atom – massive (comparatively) central, positive mass (the nucleus) with electrons moving around at a distance many times greater than the nucleus

o James Chadwick

□ Discovered the neutron

• Used the Be atom

Visual Questions

A charged particle is caused to move between two electrically charged plates. (a) Why does the path of the charged particle bend? (b) What is the sign of the electrical charge on the particle? (c) As the charge on the plates is increased, would you expect the bending to increase, decrease, or stay the same? (d) As the mass of the particle is increased while the speed of the particles remains the same, would you expect the bending to increase, decrease, or stay the same?

Millikan determined the charge on the electron by studying the static charges on oil drops falling in an electric field. A student carried out this experiment using several oil drops for her measurements and calculated the charges on the drops. She obtained the following data: (a) What is the significance of the fact that the droplets carried different charges? (b) What conclusion can the student draw from these data regarding the charge of the electron? (c) What value (and to how many significant figures) should she report for the electronic charge?

Modern View of Atomic Structure: An Introduction Section#13

• Nucleus

o Protons and neutrons (nucleus)

□ Atomic number

• Number of protons

□ mass number

• mass of nucleus

• protons + neutrons

o neutrons = mass # - protons (atomic number)

• Electron (located around nucleus at great distance)

Particle Mass Charge Location

Electron 9.11x10-31kg 1- outside

Proton 1.67x10-27kg 1+ inside

Neutron 1.67x10-27kg 0 inside

• Isotopes

o Atoms with the same number of protons but numbers of neutrons

• Atomic symbol (AZX) A – mass number, Z – atomic number, X – chemical symbol

o Sodium – 23

□ 2311Na

o Sample Exercise 2.2 – Writing the Symbols for Atoms

Write the chemical symbol for an element with atomic number of 9 and a mass number of 19? How many electrons and neutrons?

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Visual Question

Does the following drawing represent a neutral atom or an ion? Write its complete chemical symbol including mass number, atomic number, and net charge (if any).

Fill in the gaps in the table shown; assuming each column represents a neutral atom.

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Fill in the gaps in the table shown; assuming each column represents a neutral atom or an ion.

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The increments in the Celsius and Kelvin scales are equal, 1°C=1K

Matter

Heterogeneous Mixtures

Pure Substances

Elements

Compounds

Atoms

Homogeneous Mixtures

Physical methods

Physical methods

Chemical methods

Nucleus

Electrons

Neutrons

Protons

Quarks

Quarks

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