Unit 1 Chemistry Notes PDF

Unit 1 Welcome to Chemistry! chem·is·try ˈkeməstrē/ noun 1. the branch of science that deals with the identification of the substances of which matter is composed; the investigation of their properties and the ways in which they interact, combine, and change; and the use of these process...

Unit 1 Welcome to Chemistry! chem·is·try ˈkeməstrē/ noun 1. the branch of science that deals with the identification of the substances of which matter is composed; the investigation of their properties and the ways in which they interact, combine, and change; and the use of these processes to form new substances.  Unit 1 Chemical v. Physical Property Physical Property Density, luster, malleability, elasticity, melting point Chemical Property Flammability, Reactivity, nutritional information, what it will react with, etc.  Unit 1 Chemical v. Physical Change Chemical Change - A change when the beginning substance has gone through a chemical reaction (new bonds, movement of electrons, etc.). The change may be accompanied by release of heat (fire), gas (smoke), transfer of electrons (color change) Physical Change - the substance has the same chemical properties, but has taken a new form. Example, tearing a piece of paper, melting ice, balloon popping….  Unit 1 Chemical v. Physical Change Chemical or Physical Change??? 1. Metal placed in water and sinks 2. A piece of paper is torn in half 3. Food is digested 4. Iron rusts 5. Water boils  Unit 1 Notes Density A physical property of matter has a relationship between the mass (amount of substance present) and the volume (space the matter occupies). D = Mass/Volume The density of a substance is constant. The larger the mass, the more volume it will occupy This is called a direct proportion. What happens if the size of an object increases? What changes? Why do substances have various densities?  Graphing and Algebraic relationships Problem 11 on HW packet Sketch a graph. Pick an appropriate scale for the mass on the y-axis, volume on the x-axis. Mass Volume Density 15g 2cm3 __________ 60g 8cm3 __________ 120g 16cm3 __________ 480g 64cm3 __________ 750g 100cm3 __________ Calculate the density of each object. What can you say about the objects? What does the slope of the line represent?  Chapter 1 Density Problem Calculate the volume of a piece of aluminum foil that has a mass of 32.70g. When you do calculations in chemistry 1. Write the equation 2. Plug in numbers with units 3. Report answer with proper units and significant figures Variable – something in an experiment that may change from experiment to experiment or measurement to measurement. Control – Something in an experiment that will stay consistent throughout the testing process. What are some possible variables in the density lab?  Measurements in Chemistry Your measurements only as good as the tool you are using. When make measurements…… 1) Identify the units (g, kg, mL, L, etc.) 2) Identify the scale of the graduated marks (1mL , 10mL, 100mL, etc.) 3) Always report one extra decimal (a guess). The last number is always a guess between the last graduated mark shown. You can not be wrong here and we will have a variety of answers.We will see this in the next examples.  Laboratory Tools Ruler- Measures the length of an object (units are in cm)  Laboratory Tools Gram Scale- Measures the mass of an object. The last digit fluctuates (its okay! The scale is guessing)  Unit 1 Accuracy and Precision Accuracy refers to the how close you are to the actual value. Precision refers to the how close your measurements are to each other.  Section 3 Using Scientific Chapter 2 Measurements Accuracy and Precision  Section 1 Scientific Method Unit 1 Observing and Collecting Data Observing is the use of the senses to obtain information. data may be qualitative (descriptive) quantitative (numerical)  Section 3 Using Scientific Unit 1 Measurements Accuracy and Precision, continued Percentage Error Percentage error tells you how close you are to the accepted value. ** Notice there are NO absolute value bars. Some use them, so do not…..why? Measurement and Significant Figures Numbers tell a story. The numbers we use in science tell us about the measurements. What kind of equipment? What care was used? Were we just guessing? The larger the number of Sig Figs means that more care was taken in the measurement. All nonzero digits are significant. Zeros between significant figures are significant. Zeros preceding the first nonzero digit are not significant. Zeros to the right of the decimal after a nonzero digit are significant. Zeros at the end of a non decimal number may or may not be significant. (Use scientific notation.)  Significant Figures Sample Problem How many significant figures are in each of the following measurements? a. 28.6 g b. 3440 cm c. 9.10 x 10^3 m d. 0.046 04 L e. 0.006 700 0 kg Measurement and Significant Figures An exact number is a number that arises when you count items or when you define a unit. – For example, when you say you have nine coins in a bottle, you mean exactly nine. – When you say there are twelve inches in a foot, you mean exactly twelve. – Note that exact numbers have no effect on significant figures in a calculation.  Significant Figures, continued Addition or Subtraction with Significant Figures When you +/- your answer will have the same number of DECIMALS as the number with the least amount in the problem. Multiplication/Division with Significant Figures Your answer will have same number of SIG FIGS as the least number in your problem.  Section 3 Using Scientific Chapter 2 Measurements Significant Figures Sample Problem E Carry out the following calculations. Express each answer to the correct number of sig. figs. a. 5.44 m - 2.6103 m b. 2.4 g/mL × 15.82 mL  Section 3 Using Scientific Chapter 2 Measurements Conversion Factors and Significant Figures There is no uncertainty exact conversion factors. Most exact conversion factors are defined quantities. 5.4423 kg x 1000 g = 5442.3 g 1 kg  Unit 1 Scientific Notation In scientific notation, numbers are written in the form M × 10n, where the factor M is a number between 1 and 10. And n is a whole number that tells you how far the decimal moves. example: 0.000 12 mm = 1.2 × 10−4 mm Negative n means a number less than one. Positive n means a large number greater than 1.  Write these numbers in Scientific Notation. Then divide the two numbers and report with proper sig figs. 1. 1 254 000 g 2. 0.00400 cm3 Write in standard form, then multiply the numbers with proper sig figs 1. 1.443 x 105 cm 2. 9.910 x 10-3 cm  Unit 1 Scientific Notation, continued Mathematical Operations Using Scientific Notation 1. We will be using calculators. To use the calculator properly you need to know the Sci Notation button. Every calculator is different. It could be the EE, EXP, x10n, etc. example: 4.2 × 104 kg + 7.9 × 103 kg Enter the number in front first (4.2) then the sci not button. Last enter the exponent. DO NOT use the ^ or parenthesis or 10x….  Unit 1 Scientific Notation, continued Mathematical Operations Using Scientific Notation example: (5.23 × 106 µm)(7.1 × 10−2 µm) = 3.7 × 105 µm2 (do not forget to multiply the units! Warm Up 1. Calculate the following and round to proper sig figs a. 5.9932m + 2.45 x 10-1m b. 10mL x 3.1120554 g/mL c. 7.2 x 1012 km2/1.0 x 10-12km What is the length of the red line?  Section 2 Units of Measurement Chapter 2 Lesson Starter Would you be breaking the speed limit in a 40 mi/h zone if you were traveling at 60.0 km/h? one kilometer = 0.62 miles 60 km/h = 37.2 mi/h You would not be speeding! km/h and mi/h measure the same quantity using different units  Visual Concepts Chapter 2 SI (Le Systéme International d´Unités)  Section 2 Units of Measurement Chapter 2 SI Base Units  Section 2 Units of Measurement Chapter 2 Conversion Factors A conversion factor is a ratio derived from the equality between two different units that can be used to convert from one unit to the other. example: How quarters and dollars are related  Section 2 Units of Measurement Chapter 2 Conversion Factors, continued Dimensional analysis is a mathematical technique that allows you to use units to solve problems involving measurements. quantity sought = quantity given × conversion factor example: the number of quarters in 12 dollars number of quarters = 12 dollars × conversion factor  Section 2 Units of Measurement Chapter 2 SI Conversions The ending of this example can be changed to whatever unit you are using. This chart is focusing on the prefix. Our example from class was microliter( ) =  Section 2 Units of Measurement Chapter 2 Conversion Factors, continued Sample Problem B Express a mass of 5.712 grams in milligrams and in kilograms. How many Mg are in 3.44ng? The mole is one the most important numerical quantities we will use in chemistry! MOLE DAY October 23 from 6:02am – 6:02pm A mole is a quantitative value. It is a set number like……. Dozen 12 units Baker’s Dozen 13 units Ream of paper 500 sheets A fortnight 14 days A mole is a quantitative value. 1 mole is equivalent to Avogadro's number. 1 mole = 6.02 x 1023 particles This is a very, very large number! 602000000000000000000000 How big is Avogadro's number? An Avogadro's number of soft drink cans would cover the surface of the earth to a depth of over 200 miles. If you spread Avogadro's number of unpopped popcorn kernels across the USA, the entire country would be covered in popcorn to a depth of over 9 miles. If we were able to count atoms at the rate of 10 million per second, it would take about 2 billion years to count the atoms in one mole. How big is Avogadro's number? A mole of sand, spread over the United States, would produce a layer 3 inches deep. A mole of dollars could not be spent at the rate of a billion dollars a day over a trillion years. The periodic table tells us the mass of a mole for each element. 6.02 x 1023 atoms or 1 mole of…. Carbon has a mass of 12.01g Hydrogen has a mass of 1.01g Sodium has a mass 22.98g This is also called the Molar Mass (g/mol)  Section 3 Counting Atoms Chapter 3 Relating Mass to Numbers of Atoms, continued Molar Mass The mass of one mole of a pure substance is called the molar mass of that substance. Molar mass is usually written in units of g/mol. The molar mass of an element is numerically equal to the atomic mass of the element in atomic mass units.  Section 3 Counting Atoms Chapter 3 Relating Mass to Numbers of Atoms Gram/Mole Conversions Chemists use molar mass as a conversion factor in chemical calculations. For example, the molar mass of helium is 4.00 g He/mol He. To find how many grams of helium there are in two moles of helium, multiply by the molar mass.  Section 3 Counting Atoms Chapter 3 Relating Mass to Numbers of Atoms, continued Convert the following 1) 84 moles of CO2 to atoms of oxygen 2) 23 kmoles of Fe to ng challenge problem!