Chemistry Second Edition Julia Burdge
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Chemistry The Central
Copyright (c) The McGraw-Hill Companies, Inc. Permission required for reproduction or
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General Chemistry 10301 General Information: Urs Jans, MR-1218 212 650 8369
[email protected]
Your expectations: A+ understand chemistry application for real life
My expectations: - solve problems - come to class - read text book
- ask questions - take advantage of work shop
Organization of Course lab: 15% workshop/homework 5% exams: 40% (2 out of 3) final: 35% no curving, you need a C in order to take Chem10401 grade distribution Spring 2010: A: 44 (10 A+) B: 42 C: 41 D: 2 F: 8 W: 9
Syllabus schedule course outcomes office hours time
Workshop: - Workshop can be a great opportunity - Attending the workshop is mandatory - Make sure you are on the roster at the end of the first workshop!
Blackboard: ccny website current student CUNY portal
online homework - http://connect.mcgraw-hill.com/class/u_jans_cc - if no code yet, register using "free trial" (click red button "register now"; enter email address; instead of entering code click "start free trial" ) (could buy e-book (both semesters + code) for $103.35; code only $49.99)
- 9 problems for chapter 1 are posted and due on Monday
1 Chemistry: The Central Science 1.1The Study of Chemistry Chemistry You May Already Know The Scientific Method 1.2Classification of Matter States of Matter Elements Compounds Mixtures 1.3Scientific Measurements SI Base Units Mass Temperature Derived Units: Volume and Density 1.4The Properties of Matter Physical Properties Chemical Properties Extensive and Intensive Properties 1.5Uncertainty in Measurement Significant Figures Calculations with Measured Numbers Accuracy and Precision 1.6Using Units and Solving Problems Conversion Factors Dimensional Analysis Tracking Units
1.1 The Study of Chemistry Chemistry is the study of matter and the changes that matter undergoes. Matter is anything that has mass and occupies space.
The Study of Chemistry Scientists follow a set of guidelines known as the scientific method: · · · · ·
gather data via observations and experiments identify patterns or trends in the collected data summarize their findings with a law formulate a hypothesis with time a hypothesis may evolve into a theory
1.2 Classification of Matter Chemists classify matter as either a substance or a mixture of substances. A substance is a form of matter that has definite composition and distinct properties. Substances can be either elements or compounds. Substances differ from one another in composition and may be identified by appearance, smell, taste, and other properties.
Classification of Matter All substances can, in principle, exist as a solid, liquid or gas. Solids and liquids are referred to collectively as the condensed phases. The three states of matter can be interconverted without changing the chemical composition of the substance.
Classification of Matter
Solids have particles that are held closely together in an ordered fashion.
Particles in a liquid are close together but are not held rigidly in position.
Particles in a gas have significant separation from each other and move freely.
Classification of Matter An element is a substance that cannot be separated into simpler substances by chemical means. An element may consist of atoms or molecules.
Atoms of an element
Molecular form of an element
Classification of Matter A compound is a substance composed of two or more elements chemically united in fixed proportions. A compound cannot be separated into simpler substances by any physical process. Separation of a compound into its constituent elements requires a chemical reaction. A compound may consist of molecules or ions, which will be discussed later.
Molecules of a compound
Classification of Matter A mixture is a physical combination of two or more substances. A homogeneous mixture is uniform throughout. A heterogeneous mixture is not uniform throughout.
Classification of Matter A mixture can be separated by physical means into pure components without changing the identities of the components.
Classification of Matter
1.3 Scientific Measurement Properties that can be measured are called quantitative properties. A measured quantity must always include a unit. The English system has units such as the foot, gallon, pound, etc. The metric system includes units such as the meter, liter, kilogram, etc.
Scientific Measurement The revised metric system is called theInternational System of Units and was designed for universal use by scientists. There are seven SI base units
Scientific Measurement The magnitude of a unit may be tailored to a particular application using prefixes.
Scientific Measurement There are two temperature scales used in Chemistry: The Celsius scale (⁰C) Freezing point (pure water): 0⁰C Boiling point (pure water): 100⁰C The Kelvin scale (K) The absolute scale Lowest possible temperature: 0 K (absolute zero)
K = oC + 273.15
Scientific Measurement Ethylene glycol is a liquid organic compound that is used as an antifreeze in car radiators. It freezes at –12⁰C and boils at 197⁰C. Express the two temperatures using the Kelvin scale. Solution: Step 1: Use the equation
Freezing point:
–12⁰C + 273 = 261 K
Boiling point:
197⁰C + 273 = 470 K
Scientific Measurement The Fahrenheit scale is common in the United States. The size of a degree on the Fahrenheit scale is only 5/9 of a degree on the Celsius scale. Temp in oF = (9/5 x temp in oC) + 32 oF
Scientific Measurement The temperature of the surface of the sun is about 6300⁰C. What is this temperature in degrees Fahrenheit? Solution: Step 1: Use the equation
(9/5 x 6300⁰C) + 32⁰F = 1.1x104 ⁰F
Scientific Measurement There are many units (such as volume) that require units not included in the base SI units. The derived SI unit for volume is the meter cubed (m3). A more practical unit for volume is the liter (L). 1 dm3 = 1 L 1 cm3 = 1 mL
Scientific Measurement The density of a substance is the ratio of mass to volume.
d = density m = mass V = volume SI-derived unit: kilogram per cubic meter (kg/m3) Other common units: g/cm3 g/mL g/L
(solids) (liquids) (gases)
Scientific Measurement Given that 25.0 mL of mercury has a mass of 340.0 g, calculate (a) the density of mercury and (b) the mass of 120.0 mL of mercury.
Solution: Step 1: Use the equation d = m/V to calculate density or mass as needed. Part (a) d = m/V = 340.0 g / 25.0 mL = 13.6 g/mL Part (b) 3
m = d x V = 13.6 g/mL x 120.0 mL = 1.63 x 10 g
1.4
The Properties of Matter
There are two general types of properties of matter: 1) Quantitative properties are measured and expressed with a number 2) Qualitative properties do not require measurement and are usually based on observation.
The Properties of Matter A physical property is one that can be observed and measured without changing the identity of the substance. Examples: color, melting point, boiling point
A physical change is one in which the state of matter changes, but the identity of the matter does not change. Examples: changes of state (evaporation, melting, condensation, freezing, ...)
The Properties of Matter A chemical property is one a substance exhibits as it interacts with another substance. Examples: flammability, corrosiveness A chemical change is one that results in a change of composition; the original substances no longer exist. Examples: digestion, combustion, oxidation
The Properties of Matter An extensive property depends on the amount of matter. Examples: mass, volume An intensive property does not depend on the amount of matter. Examples:
temperature, density
1.5 Uncertainty in Measurement There are two types of numbers used in chemistry: 1) Exact numbers which have defined values 1 kg = 1000 g 1 dozen = 12 object any number obtained by counting 2) Inexact numbers numbers obtained by any method other than counting
Uncertainty in Measurement An inexact number must be reported so as to indicate its uncertainty. Significant figures are the meaningful digits in a reported number. The last digit in a measured number is referred to as the uncertain digit.
Uncertainty in Measurement The number of significant figures can be determined using the following guidelines: 1)
Any nonzero digit is significant 1129 m
2)
4 significant figures Zeros between nonzero digits are significant 109 cm 3 significant figures
3)
Zeros to the left of the first nonzero digit are not significant
0.0003 kg
1 significant figure
Uncertainty in Measurement The number of significant figures can be determined using the following guidelines: 4) Zeros to the right of the last nonzero digit are significant if a decimal is present. 9.550 m
4 significant figures
5) Zeros to the right of the last nonzero digit in a number that does not contain a decimal point may or may not be significant. 1200 m
ambiguous
Uncertainty in Measurement In addition and subtraction, the answer cannot have more digits to the right of the decimal point than any of the original numbers. 102.50 ← two digits after the decimal point + 0.231← three digits after the decimal point 102.731 ← round to 102.73 143.29 ← two digits after the decimal –20.1 ← one digit after the decimal 123.19 ← round to 123.2
Uncertainty in Measurement In multiplication and division, the number of significant figures in the final product or quotient is determined by the original number that has the smallest number of significant figures. 1.4 x 8.011 = 11.2154 figures)
←
round to 11 (limited by 1.4 to two significant
11.57/305.88 = 0.0378252 ← round to 0.03783 (limited by 11.57 to four significant figures)
1.46: Express the answer to the following calculation in scientific notation: a) 147.75 + (2.3 x 10-1)
b) 79,000 : (2.5 x 102)
Uncertainty in Measurement Exact numbers can be considered to have an infinite number of significant figures and do not limit the number of significant figures in a result. In calculations with multiple steps, round at the end of the calculation to reduce any rounding errors.
Uncertainty in Measurement An empty container with a volume of 150.0 cm3 is weighed and found to have a mass of 72.5 g. The container is filled with a liquid and reweighed. The mass of the container and the liquid is 194.37 g. Determine the density of the liquid to the appropriate number of significant figures. Solution: Step 1: Determine the mass of the liquid 194.37 g – 72.5 g 121.87 g
← do not round this number yet, but remember it is good to only one decimal place or a total of 4 sig. figs.
Step 2: Calculate the density d = 121.87 g / 150.0 mL = 0.812466 g/mL ← this number rounds to 0.8125 g/mL
Uncertainty in Measurement Accuracy tells us how close a measurement is to the true value. Precision tells us how closely multiple measurements of the same thing are to one another.
Good accuracy and good precision Poor accuracy but good precision Poor accuracy and poor precision
Uncertainty in Measurement Three students were asked to find the mass of an aspirin tablet. The true mass of the tablet is 0.370 g.
Student A: Results are precise, but not accurate Student B: Results are neither precise nor accurate Student C: Results are both precise and accurate
Uncertainty in Measurement A conversion factor is a fraction in which the same quantity is expressed one way in the numerator and another way in the denominator. 1 in = 2.54 cm
Uncertainty in Measurement The use of conversion factors in problem solving is called dimensional analysis or the factor-label method. Example: Convert 12.00 inches to meters. 1 in = 2.54 cm (exact) 1 m = 100 cm
Uncertainty in Measurement The American Heart Association recommends that healthy adults limit dietary cholesterol to no more than 300.0 mg per day. Convert this mass of cholesterol to ounces. 1 oz = 28.3459 g Solution: Step 1: The problem requires a two step dimensional analysis conversion: mg to grams; grams to oz.
1.56: The average speed of helium at 25 oC is 1255 m/s. Convert this speed to miles per hour (mph). 1 mi = 1609 m
Online Homework Problems for Chapter 1: 16, 24, 38, 46, 48, 50, 56, 63, 106
