Ch. 1 Introduction: Some Basic Concepts
-
Upload
clare-farley -
Category
Documents
-
view
29 -
download
2
description
Transcript of Ch. 1 Introduction: Some Basic Concepts
Ch. 1 Introduction: Some Basic Concepts
Chemistry
• The study of the composition of matter and the changes that matter undergoes
Matter
• Anything that has mass and takes up space
States of Matter
Solid
• Definite shape AND volume• Incompressible• Particles are packed tightly together
Liquid
• Definite volume• Takes the shape of its container• Almost incompressible-particles not rigidly
packed
Gas• Takes the volume and the shape of its
container• Particles in a gas are spaced far apart• Easily compressed
Vapor
• The gaseous state of a substance that is generally a solid or a liquid at room temperature
Physical Properties• A quality or condition that can be observed or
measured without changing the substances composition
• Color, solubility, odor, hardness, density, melting point, boiling point
Physical Change
• Changes the material without changing the composition
• Boiling, freezing, dissolving, melting, condensing, breaking, splitting, cracking, cutting, crushing, bending………
• Usually reversible
Pure Substance
• Contain only one kind of matter• Have identical physical properties
Elements
• The simplest forms of matter that can exist under laboratory condition
• Can not be separated into simpler substances by chemical means
• The building blocks for all other substances
Chemical Symbol
• One or 2 letters• The first letter is always capitalized
Compounds
• 2 or more elements chemically combined• Can be separated into simpler substances by
chemical means
Sodium Metal
+Chlorine Gas
=Sodium Chloride
Compounds• The subscript numbers in chemical formulas
represent the proportions of elements that make up the compounds
• Pb2(SO4)3
• Pb (lead) 2• S(Sulfur) 3• Oxygen 12
Law of Definite Proportions (Law of Constant Composition)
• The elemental composition of a pure compound is ALWAYS the same
• i.e. Water= H2O = 2 Hydrogens: 1 Oxygen ALWAYS
2 Types of Mixtures
• Heterogeneous Mixtures• Homogeneous Mixtures
Mixture
• A physical blend of 2 substances• Compositions may vary
Heterogeneous Mixture
• One that is NOT uniform in composition
Homogeneous Mixture
• The same throughout
Solution
• A homogeneous mixture• Solutions may be solids, liquids, or gases• Same composition throughout
Phase
• Any part of a system with uniform composition and properties
Phase 1
Phase 2
Separation of Mixtures
Distillation
Separation of Mixtures
Centrifuge
Separation of Mixtures
Chemical Property• The ability of a substance to undergo a
chemical reaction to form a new substance
• I.e. flammability, alkalinity, acidity, rusting
• Chemical properties are only observed when a substance undergoes a chemical change
Chemical Changes (Reaction)
• One or more substances change into new substances
• 2 H2 (g) + O2 (g) 2 H2O (g)
Reactants Products
Chemical Reactions (rust-oxidation)
Chemical Reactions (combustion)
Chemical Reactions (acid/base)
Indicators of a Chemical Reaction
1. Energy is absorbed or given off (change in temperature)
2. Change in color3. Change in odor4. Formation of a solid (precipitation)5. Formation of a gas
Scientific Method
• One logical, systematic approach to the solution of scientific problems. Steps include:
1. Making observations2. Testing hypothesis3. Developing theories
Observation
• Use your senses to obtain information directly
Hypothesis
• A proposed explanation for an observation based on previous knowledge (or research)
• Must be specific• Must be testable• Is only useful if it accounts for what is actually
observed
Experiment
• A means to test a hypothesis
Manipulated Variable (Independent Variable)
• The variable that you can change• Time• Temperature• Volume• Speed• Pressure
Independent Variable
Responding Variable (Dependent Variable)
• The variable that is observed during the experiment
Dependent Variable
• For the results of an experiment to be accepted the experiment must produce the same results no matter how many times it is repeated or by whom
Theory
• A broad and extensively tested explanation of why experiments give certain results.
• A theory can NEVER be proven because a new experiment can always disprove it
Scientific Law
• A concise statement that summarizes the results of many observations and experiments.
• Scientific law describes natural phenomena without attempting to explain it
Qualitative Observation
• Give results in a descriptive non-numerical form
• Subjective• Ex. The solution is green, The precipitate is
fluffy
Quantitative Observation
• Gives results in a definite form• Numbers and units
Measurement
• A quantity that has both a number and a unit
• Measurements are only as exact as the instrument used to take it
International System of Units (SI)
• Length = meter(m)• Mass = gram (g)• Temperature = kelvin (K) although often we
will use Celsius• Time = second (s)• Amount of substance = mole (mol)• Volume = liter (L)
Prefixes
• Mega (M) = 106
• Kilo (k) = 103
• Deci (d) = 10-1
• Centi (c) = 10-2
• Milli (m) = 10-3
• Micro () 10-6
• Nano (n) = 10-9
• Pico (p) = 10-12
Length
• The distance between 2 points• Unit: meter (m)
Mass
• The amount of matter an object has• Units: Grams• Measuring tools: Triple Beam Balance,
Electronic scale, Analytical Scale
Temperature
• Measure of how hot or cold an object is• Determines heat transfer (moves from high to
low)• Almost all substances expand when heated-
contract when cooled (except water)• Units: Kelvin (K) = Celsius (c) + 273
Volume
• The amount of space an object (substance) occupies
• Units: Liter (L)
Density
• The amount of matter in a given volume
• Density = mass volume• D= m/v
Density
• Depends only on the composition of a substance NOT on the size of the sample
• With a mixture the density can vary because the composition of the mixture can vary
Density
• The density of a substance generally decreases as its temperature increases therefore temperature of the substance must always be noted
Accuracy
• how close a measurement comes to the actual or true value of whatever is measured must be compared to the correct value
Precision
• How close a series of measurements are to one another depends on more than one measurement
Significant Figures
• In a measurement, includes ALL of the digits that are known plus the last digit that is estimated
• Measurements MUST ALWAYS be reported to the correct number of sigfigs because calculated answers depend on sigfigs
Rules for Significant Figures
1. Nonzero numbers are ALWAYS significant-13.24 cm = 4 sigfigs- 3.5 mL = 2 sigfigs-123.456 g = 6 sigfigs
Rules for Significant Figures
2. Captive zeros are ALWAYS significant- 1023 g = 4 sigfigs- 1.0005 cm = 5 sigfigs
Rules for Significant Figures
3. Leading Zeros (zeros in front of nonzero numbers) NEVER count- 0.53 cm = 2 sigfigs- 0.0000000000053 cm = 2 sigfigs
Rules for Significant Figures
4. Zeros at the end of a number only count if there is a decimal point- 100 mL = 1 sigfig-100.0 mL = 4 sigfigs-100.00 mL = 5 sigfigs
Rules for Significant Figures
5. Scientific Notation: Only consider the coefficients when determining sigfigs- 1.2 x 104 mL = 2 sigfigs- 1.06 x 10-6 m = 3 sigfigs
Rules for Significant Figures
6. Exact numbers (counting) or exactly defined quantities have unlimited number of sigfigs- 35 students = unlimited sigfigs- 1 min = 60 seconds = unlimited sigfigs
Significant Figures in Calculations
• A calculated value can NEVER be more precise than a measured value
Significant Figures in Calculations
• Addition and subtraction:• The answer should be rounded to the same number of
decimal places as the measurement with the least number of decimal places
• 135.75 mL (2 decimal places)+ 57.6 mL (1 decimal place) 193.35 mL = 193.4 mL (answer has 1 decimal place)
Rounding
• X > 5 Round up
• X < 5 Round down
• Only look at the number next to the one you are rounding
Significant Figures in Calculations
• Multiplication and Division• Round the answer to the same number of significant
figures as the number with the least number of sigfigs.
• 35.062 g ( 5 sigfigs) = 1.492 g/mL (4 sigfigs) 23.50 mL (4 sigfigs)
Dimensional Analysis
1. Always start with the number given in the problem)
2. Units that you want to cancel go on the bottom
3. When a number is on top, multiply4. When the number is on the bottom, divide
Example
• Convert 153 lb to grams (hint: 1 lb = 453.6 g)
Example
• 153 lb X 453.6 g = 69400.8g=69400g (3sigfigs) 1 lb