• You must learn about matter on three levels, macroscopic, symbolic and atomic.
• Practice Exercise: describe or sketch the substance water in the following ways:
  • Macroscopic view:
  • Symbolic view:
  • Atomic view:

• It is a pure substance that cannot be broken down into simpler substances.
• It is a large collection of one type of particle.
• It is represented by a chemical symbol which uses 1 or 2 letters of the elements name with the first letter capitalized.
• It may exist as a solid, liquid or gas.
• It may exist in various aggregates of two or more atoms.
• It may be reactive or unreactive.
• Examples: Figure 2.1, Figure 2.9, Figure 3.1

• Use 1 or 2 letter abbreviations
• Capitalize the first letter only
• Some symbols come from Latin names
• Examples:
  • C carbon
  • Cu copper (cuprum)
  • N nitrogen
  • Au gold (aurum)
  • F fluorine
  • Na sodium (natrium)

• Concepts are a single word or phrase that represents a complex idea. (These are often given in bold type in your text.)
  • A “definition” contains the essential characteristics of a concept.
  • Variable characteristics: not all concepts have these characteristics, but most do.
• Examples and non examples help clarify concepts and their relationships.

• Practice the definition by writing it in your own words using its essential characteristics.
  • Write your own definition of the concept: element.
• Determine its variable characteristics.
  • List the variable characteristics of the concept: element
• Try to find as many examples and non-examples as you can.
  • Identify one example of an element and explain which essential and variable characteristics are present or not present.
• Use outlines, pictures or diagrams to understand the characteristics of a concept.
  • Are the pictures of an element in Figure 2.1, Figure 2.9, Figure 3.1 macroscopic, symbolic or atomic?
• Increase your sensitivity to the exact meaning of chemical concepts and develop clear relationships to related concepts
  • What is the difference between an element and an atom?

• The smallest unit of an element that retains the microscopic properties of the element.
• It can exist alone or combined.
• It may be in any physical state.
• An atom can be represented by a sphere or other generic shape.

• Many different concepts are taught in a single class.
• Concepts are not always presented or without explicit links to related concepts.
• They are taught simultaneously with other types of content.
• Concepts are extended in later class periods.

• Multiple Choice: identify a concept word, a definition, essential characteristics or examples
• Short Answer: supply a definition for a concept and/or give an example
• Matching: link a concept to an example, discriminate between related concepts
• T/F: identify correct examples or characteristics

• All atoms of an element are identical in mass and properties.
• Compounds are formed by combinations of different kinds of atoms.
• Atoms combine in small, whole number ratios.
• Atoms are the units of chemical change but are not created, destroyed or converted into other atoms during reactions.

• Combining Dalton’s ideas about atom with our current understanding of this concept, has created the “particulate model of matter”:
  • Atoms are often symbolically represented as spheres.
  • Atoms of the same element have the same color, size or shading, which is NOT necessarily related to the properties of the atom.
  • When atoms combine to form compounds, the atomic spheres are attached to each other in some way.
  • The number of spheres of each type of element represents the proportion of elements in that substance.

• Pure Substance - fixed composition but only one chemical substance present; the most common examples are elements and compounds
  • Element - one type of atom present; cannot be decomposed by chemical reactions
  • Compound - specific ratio of two or more atoms that are joined together
• Heterogeneous Mixture - nonuniform, variable composition of two or more substances such as whole blood or orange juice
• Homogeneous Mixture - uniform but variable composition with more than one substance present; the most common example is a solution such as blood plasma or brewed coffee
• A mixture or pure substance can be in any physical state
  • solid - not fluid or compressible, has a defined shape and volume; particles are not moving freely through the container and are packed neatly together
  • liquid - fluid but only slightly compressible, assumes the shape of a container but does not always fill the entire volume; particles are moving freely in part of the container and are close to each other
  • gas - fluid, very compressible, assume the shape of a container and fills its volume; particles are moving freely and randomly and are far apart

Chemical vs. Physical Changes

• Physical changes: when two or more substance are combined or separated but, each substance retains its chemical identity
  • Describe how these characteristics are illustrated when CuSO₄ is added to water
• Physical changes: changes in physical state
  • Describe how these characteristics are illustrated when ice melts to water.
• Chemical changes: when a chemical substance forms a new combination of atoms
  • Describe how these characteristics are illustrated when CuSO₄ is added to NaOH.
• Indicators of physical changes are changes
  • in density
  • in hardness
  • in texture
  • in phase (solid, liquid, gas)
• Indicators of chemical changes are:
  • a color change
  • formation of a solid precipitate
  • formation of a gas
  • a spontaneous increase or decrease in temperature

• Every measurement is a number followed by the measuring unit used.
• You are making a measurement when you check your weight, read your watch, take your body temperature, etc.
• In science the metric system is used for measurements.
  • It is a decimal system based on a unit of 10.
  • Prefixes increase or decrease the magnitude of a number by 10’s.

• Learn the facts in Table 1.2 (SI Base Units); skip luminous intensity.
• Other important derived units: Volume (liter = L), Energy (joule = j and calorie = cal) temperature (Celsius = °C).

• Prefixes increase or decrease the basic unit by 10’s
• Example: 1 kilometer = 1 km = 1000 meters
• Commonly used prefixes that increase a unit are: giga (G = 1,000,000,000), mega (M = 1,000,000), kilo (k = 1000)
• Commonly used prefixes that decrease a unit are: deci (d = 0.1), centi (c = 0.01), milli (m = 0.001), micro (µ = 0.000001), nano (n = 0.000000001), pico (p = 0.000000000001)

• Heat is an intensive property of matter and represents the transfer of thermal energy which is based on the random motion of molecules
  • Heat is measured using: Celsius (°C), Kelvin (K) and Fahrenheit (°F).
  • 1 °C (Celsius scale) = 1 K (Kelvin scale); K = °C + 273
  • °C = (°F -32)* 5/9; °F = (°C * 9/5) + 32
• Energy is an extensive property and represents the ability of a substance to do work.
  • Energy is measured in calories or joules
  • calorie = heat required to raise 1 g water by 1 degree C.
  • 1.00 calorie = 4.184 J = heat capacity of water in cal/g*°C
  • 1 kcal (kilocalorie) = 1 Calorie; kJ = kilojoule = 1000 J

• Facts are true statements: names, descriptions of properties measured values, etc.
• Example: Water is a colorless liquid that freezes at 0.0^o C.
• Facts can only be memorized!

• You have a short and long term memory.
• Information is first placed in the short term memory where it is is lost quickly.
• “Memorization” for learning is actually transfer from short to long term memory.
• In both types of memory the links fade over time and must be reinforced.

• Use as many senses as possible to help; read, write and speak the information out loud!
• Look for patterns that will help organize what you have to memorize.
• Group the items into chunks that will fit into your short term memory.
• Practice the items in random order.
• Spread the memorization over time. “Cramming” fills only the short term memory.

• Straight Recall
  • What is the symbol for milliliters?
• Recognition of Facts and Errors
  • Which is the smallest SI unit?
    • (a) deci (b) centi (c) kilo (d) pico
  • The symbol for milliliter is mL. T or F?
  • Use of memorized information in other questions and problems
    • Convert the quantity 10.0 mL to liters.

• It uses a coefficient and power of 10 to represent a decimal number.
  • Example: 0.000001 = 1x10^-6; 1000 = 1x10³.
  • Non examples: 1, 1/2, 0.5, 10^0.5
• The power of ten factor indicates how many places the decimal has been moved; right = negative, left = positive.
• Scientific notation can usually be activated on your calculator with the EE or EXP function key. (Do NOT type in: 1 times 10 minus 6 for 1x10^-6.)

• Using a measuring tool to determine a quantity.
• Examples: your height, weight or temperature.
• Measured numbers have the variable characteristics of accuracy and precision.
• Different levels of precision occur when measured numbers are estimated between calibration marks. The number reported includes the last estimated value and this determines the number of “significant figures”.

• Integers obtained by counting such as: 2 soccer balls, 1 watch, 4 pizzas
• Values obtained from unit equivalencies: 1 foot = 12 inches, 1 meter = 100 cm.
• Exact numbers have no estimated digits and are exempt from “significant figure” rules.
• Both measured and exact numbers can be written using scientific notation.

• In the number 2.76 the known digits are 2 and 7 and are 100% certain.
• The third digit, 6, is estimated and uncertain.
• In reporting this value all measured and estimated digits are “significant”.

• Precision
  • For single measurements, the calibration of a measuring device and how it is used determines the precision based on the number of significant figures.
  • For multiple measurements precision is a measure of how well several determinations of the same quantity agree with each other
  • Example: experimental values for the density of Mg were 1.685 g/mL, 1.69 g/mL, 1.67 g/mL, 1.7 g/mL
  • The level of precision is fixed by the technique used in making the measurements. For example, using a beaker versus a buret to measure fifty milliliters of solution.
• Accuracy
  • how well the measured quantity agrees with a reference or accepted value.
  • example: accepted density of Mg is 1.74 g/mL, the measured density is 1.7 g/mL.
  • The level of accuracy is determined by the situation. For example, hitting the edge of an archery target from 10 feet versus 1000 feet.

• SF in a measurement includes the all the known digits and one estimated digit.
• Determine the number of SF in each measured or given value by counting the number of digits left to right starting with the first nonzero digit.
• Examples: 38.15 cm (4 SF), 5.6 ft (2 SF), 65.6 lb (? SF), 122.55 m (? SF)
• Examples with Leading Zeros: 0.008 mm (1 SF), 0.0156 oz. (3 SF), 0.0042 lb. (? SF), 0.000262 mL (? SF)
• Examples with Sandwiched Zeros: 50.8 mm (3 SF), 2001 min (4 SF), 0.702 lb. (? SF), 0.00405 mL (? SF)
• Examples with Trailing Zeros: 25,000 in (2 SF), 200 yr. (1 SF), 48,600. gal. (5 SF), 25,005,000 g (? SF)

• A calculated answer must match the least precise measurement in the number of SF. Example: 1.1/1.1 = 1 ---> 1.0 since there are 2 SF in each measured value.
• To round a number off to the proper number of significant figures: the last digit retained is increased by 1 only if the following digit is 5 or greater.
• You may apply the rules for rounding calculated numbers either at each major step in a calculation chain or at the end only.
• For addition and subtraction: the answer has the same number of decimal places as the measurement with the fewest decimal places. Example: 9.2 + 1.34 = 10.5 (the number 9.2 is the value with the fewest decimal places so we use that many decimal places in the answer)
• For multiplication and division: round (or add zeros) to the answer until you have the same number of significant figures as the measurement with the fewest total significant figures. Example: 1.34 / 25.20 = 0.053174603 is rounded to 3 SF or 0.0532.

• Rules free you from learning many related individual facts but must be applied under specific circumstances.
• A rule is any generalization that summarizes chemical behavior. Examples:
  • Laws:
  • Chemical equations:
  • Math formulas and graphs:
  • Explicit rule sets:
  • Generalizations of chemical facts:

• Learn the underlying facts and concepts.
  • What concepts are directly related to SF?
• Restate the rule in your own words.
  • Take one set of examples for determining the SF in a measured value and create a rule statement for it.
• Identify when a rule is applied and what the rule accomplishes.
  • When are the rules of SF to be applied in this course?
• Identify exceptions to the rule.
  • What types of numbers are exempt from SF?
• Make diagrams or flow charts to illustrate rules.
  • Create a flow chart that shows how to determine the SF in a number.

• Some rules must be memorized, particularly those that cannot be “figured out” and seem arbitrary.
• Link rules to concepts in as many ways as possible. This will provide a context or framework to help you learn the rule and apply it correctly.
• Make sure the rules give the same answer all the time; continually check the validity and application of the rule.
• Develop your own chemical rules particularly when faced with tables or graphs.

• Questions which ask you to apply a rule are the most common type on exams.
  • Short Answer: What is the temperature of the boiling point of water in degrees K?
  • Multiple Choice: The number of significant figures in 0.0060 is (a) 4 (b) 3 (c) 2
• Problems in Applying Rules
  • Questions that test exceptions to rules: Are the zeros in the number 0.00044 significant?
  • Choosing among alternative rules: Give the decimal equivalent to the following expression to the correct number of significant figures: (0.0043 + 1.2345)/.554.
  • Combined rules: What is the temperature of the boiling point of water in degrees K and degrees F?

• Equalities can be used to create conversion factors; example: 1 m = 100 cm can be converted to 1m/100 cm or 100 cm/m.
• Focus on metric-metric conversion factors, measured number to exact number conversion factors and percent conversion factors.
• Examples: 100 cm/m, 10.0 mg/1 tablet, 20.0% = 0.200
• These conversion factors can be used to solve problems using generic techniques called “dimensional analysis” or the “factor-label” method.

• Use the steps below to answer the following example problem: A physician ordered 1.0 g of tetracycline to be given every six hours to a patient. If your pharmacy has only 500. mg tablets, how many tablets will you need for 1 day’s treatment?
  • Identify or summarize all the numerical data given including their unit labels.
  • Identify the answer you wish to obtain and its units.
  • Develop a unit conversion plan to change one unit into another.
  • Combine (usually multiply) the conversion factors so that each unit that does not appear in the unit of the final answer “cancels out” by appearing somewhere in both the numerator and denominator of different factors.
  • Calculate the final answer and check its validity.

• Which weighs more a kilogram of rocks or feathers?
• Which weighs more a liter of rocks or feathers?
• Will rocks or feathers float on water? On mercury? Why?
• Density = mass/volume
• Typical units for this measurement: (g/mL) or g/cm³.
• Generic problems involving density are best solved using the mathematical formula method.

• Use the steps below to answer the following example problem: A thermometer contains 8.3 g of mercury and the density of mercury = 13.6 g/mL. What volume of mercury is present?
  • Identify the mathematical relationship that relates or includes all the given data.
  • Identify the unknown quantity and assign numerical values to the known quantities.
  • Substitute the known values in the mathematical relationship.
  • Algebraically transform the equation so that the unknown is alone on one side.
  • Calculate the final answer and check its validity.

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