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Chemistry: A Molecular Approach

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1 Chemistry: A Molecular Approach
The properties of matter are determined by the properties of atoms and molecules. The understanding of matter at the molecular level gives us the ability to control that matter. First semester General Chemistry – the primary goal is to learn the fundamental principles of Chemistry.

2 Why are you here? Why study chemistry? Reasons 1) 2) 3)

3 Ch. 1: Matter, Measurement…
A comparison of two molecules: CO consists of one carbon and one oxygen atom. CO2 consists of one carbon and two oxygen atoms. One is a poison while the other is harmless.

4 Ch. 1: Matter, Measurement…
Matter can exist in three different states: Solid – Is rigid, has both a definite shape and volume. Solids can by crystalline or amorphous. Liquid – Is a fluid – molecules or atoms can flow. However, it has a fixed volume but no fixed shape. Gas – Also a fluid. However, it has no fixed volume or shape. Gases are unique in that they can be compressed. Liquids and Solids cannot be compressed. Molecular View

5 Learning Check What phase(s) can be compressed?
What phase(s) have a fixed volume? What phase(s) have a fixed shape?

6 Classification of Matter
Pure substance – has a fixed composition and distinct properties. A pure substance can be either an element or a compound. Ex) Table sugar (sucrose) and Carbon are pure substances.

7 Classification of Matter
Elements are substances that cannot be decomposed into simpler substances. Each element is composed of only one kind of atom. Compounds are composed of two or more elements that have been chemically combined. Ex) Sulfur and Sodium chloride

8 Classification of Matter
Most matter we encounter is a mixture of substances. The elements and / or compounds in a mixture always retain their identity. Ex) Gatorade and Milk

9 Classification of Matter
Homogenous – the mixture is uniform throughout. Most homogenous mixtures are solutions. Crystal Light is a solution. Heterogeneous – the mixture is not uniform throughout. A chocolate chip cookie is a heterogeneous mixture.

10 Separation of Mixtures
Matter Variable composition? No. Pure Substance Separation by ____________? Element Yes. Compound Mixture ________________? Heterogeneous Homogeneous

11 Learning Check Decide whether the following is a mixture or a pure substance. Tomato juice Iodine crystals Sand Baking soda A substance does not have a variable composition and can be separated by chemical means. This is a _________________________.

12 Separating Mixtures Separating mixtures into their components is something that scientists often do. Simple: filtering sand from water. Simple: using a magnetic to collect iron fillings from sand. Hard: Distillation of two liquids. Hard: Chromatography.

13 Properties of Matter Physical property ‑ a characteristic that can be observed for a material without changing its chemical identity. extensive property ‑ dependent on the amount ex) mass, volume, heat content intensive property ‑ independent of the amount ex) density, temperature, melting point

14 Physical Properties Copper Reddish, shiny Melting point = 1085oC
Density = 8.94 g/mL Specific heat = J/g C Crystal structure = face centered cubic

15 Properties of Matter Chemical property - describes how a substance may change or react towards other substances. Ex) Propane burns in air to form carbon dioxide and water

16 Physical and Chemical Changes
Physical change ‑ a change in the form of matter but not it’s identity ‑ can often easily return to the former state ex) melting of ice ex) dissolving of salt in water ex) ripping a piece of paper

17 Physical and Chemical Changes
Chemical change ‑ a change in which one or more kinds of matter are transformed into new kinds of matter ‑ difficult to return to the former state (by any physical means) ex) rusting of iron ex) burning of wood

18 Learning Check Would classify these common daily activities as a physical or chemical change? Getting a haircut Applying bleach to turn your hair blond Baking a cake Adding a Crystal Light© packet to water Leaves on a tree turning red or yellow Rubbing alcohol evaporating from your skin Pouring spaghetti sauce over pasta

19 Energy Most physical and chemical changes involve changes in energy.
Example – water evaporating from your skin. Example – Burning propane in an outdoor grill. Energy = capacity to do work. Work = Force times distance.

20 Types of Energy Potential Energy = energy of position or composition.
Kinetic Energy = energy of motion. Law of conservation of energy. Tendency of systems with high PE.

21 Units of Measurement Units used in scientific measurements are those of the metric system. Although we still use the English system, the metric system is becoming more common.

22 SI units 1960 agreement on a set of internationally accepted group of seven base units from which all others are derived. Mass = kilogram Length = meter Time = second Temperature = Kelvin Amount of Substance = mole Electric current = Ampere Luminous intensity = Candela

23 Metric Prefixes Common ones include: Non standard unit, Angstrom’s (Å)
Kilo (k) = 103 Centi (c) = 10-2 Milli (m) = 10-3 Micro (m) = 10-6 Nano (n) = 10-9 Non standard unit, Angstrom’s (Å) 1 Å = 1 x 10-8 cm

24 Temperature Scales Temperature is a measurement of the hotness or coldness of an object. English scale = Fahrenheit Celsius scale – assigns temperatures based on melting and boiling points of water. Kelvin scale – based on absolute zero being the coldest possible temperature K = oC

25 Derived Units Examples are units like volume, density, and velocity.
Volume of a cube = length cubed Density of a substance = mass divided by volume.

26 Uncertainty in Measurement
Exact numbers – a number that is known to be precisely that value. These have no effect on sig. figs. Ex) 12 inches = 1 foot Ex) 15 apples in a bag Inexact (measured) numbers – have some amount of uncertainty. Ex) A coin has a mass of 2.52g Ex) A bottle of soda has a volume of 591mL

27 Uncertainty in Measurement
Any measurement contains some uncertainty. precision ‑ the closeness of a group of figures to each other ‑ standard deviation accuracy ‑ the closeness of a single value or an average to the accepted value - percent error Density of Al Student A Student B Student C Trial 1 2.5 g/mL 2.58 g/mL 2.68 g/mL Trial 2 3.3 g/mL 2.51 g/mL 2.72 g/mL Trial 3 3.1 g/mL 2.55 g/mL 2.69 g/mL Average 3.0 g/mL 2.70 g/mL %error / %rad

28 Study Check Do the following represent Exact or Measured numbers?
A store has 35 bicycles on display. The density of an object is found to be 1.8g/mL. 1 meter is equivalent to 100 centimeters. Planck’s constant is listed in the book as x 10-34J s. There are about 454 grams in one pound.

29 Significant Figures Significant figures are the digits measured in a number such that all certain digits plus one uncertain digit is included. Certain digits – all performing measurement would agree on these. Uncertain digit – a “best guess” – it is each individual’s best interpretation of the measurement.

30 “Best Guess” examples

31 Rules: Non-zero numbers are always significant.
Zeros between non-zero numbers are always significant. Zeros at the beginning of a number are NEVER significant; they are merely placeholders. Zeros that fall at the end of the number and after the decimal point are always significant. When a number ends in zeros without a decimal point, the zeros may or may not be significant. We will err on the least number.

32 Scientific Notation Puts all numbers in the form of: A x 10n, where A is a number between 1 and 10 and n is the exponent equally to the number of places the decimal point must be moved. Scientific Notation removes ALL ambiguity from determining significant figures. See Appendix A for a review of scientific notation.

33 Examples How many significant figures do these measured numbers have?
0.0092 9.20 92,000 92,000.0

34 Sig. Figs. In Calculations
Multiplication and Division ‑ your answer will keep only the same number of sig. figs. as the measurement that had the fewest number of sig. figs. Addition and Subtraction ‑ your answer will have the same number of sig. figs. as the one with the fewest decimal places. Note – in series of calculations, do not round until the very end. In mixed calculations, follow rules for each individual calculation.

35 Rounding of Numbers Rounding is the process of dropping non-significant digits in a calculation and adjusting the last digit. Rules: if the leftmost digit (to be dropped) is 5 followed with any non-zero digits, then round the final digit up one. if the leftmost digit is less than 5, round down. if the digit is 5 only, and the digit to be rounded is even, round down. if it is odd, round up.

36 Examples Round each calculation to the correct number of significant figures. 1.305 x = 105.2 x = 495.0 ÷ 0.23 =

37 Examples Round each calculation to the correct number of significant figures. = 27.86 – = 1.8 121 – =

38 Examples ( – ) x = 2.5 ( ) ÷ = (48.32 – 4.5) ÷ =

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