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Chem 121

Chem 121. Introduction to Inorganic Chemistry. What is Matter?. Matter is anything that has mass and occupies space. Mass is a measurement of the amount of matter present. The mass is constant, no matter where it is …on the moon or on earth.

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Chem 121

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  1. Chem 121 Introduction to Inorganic Chemistry

  2. What is Matter? • Matter is anything that has mass and occupies space. • Mass is a measurement of the amount of matter present. The mass is constant, no matter where it is …on the moon or on earth. • Weight the a measurement of the gravitational force pulling the object toward earth. The weight of an object will be different on the moon that has a different grav. pull.

  3. Properties and Changes • Two types of properties: • Physical : Those that can be observed or measured without changing or trying to change the composition of the matter in question * no original substances are destroyed and no new substances are created. Like Color and SIZE. • Chemical: properties that matter demonstrates when attmepts are made to change it into other kinds of matter.

  4. Physical Change VS Chemical Change • Physical Change: Cutting a piece of paper, Boiling water, freezing water, heat is added or removed from matter. • Chemical Change: burn the paper. ,mix the water with an acid.

  5. A Model of Matter • Scientific Models are explanations for observed behavior. • All matter is made up of particles that are too small to see. (molecules) • Molecules: The smallest particle of a pure substance that has the properties of that substance and is capable of a stable independent existence. A molecule is also the limit of physical subdivision for a pure substance.

  6. Example • Oxygen: helps a substance burn more rapidly..like wood. • A large amt. or a small amt of Oxygen would behave the same. • The smallest amt. that would still behave the same is known as the molecule.

  7. What’s beyond the Molecule? • John Dalton wanted to know. In 1808 he proposed the following: • 1. All matter is made up of tiny particles called atoms. • 2. Substances called elemenats are made up of atoms that are all identical. • 3. Substances called compounds are comginations of atoms or two or more elements.

  8. Cont. • 4. Every molecule of a specific compound always contains the same number of atoms of each kind of element found in the compound. • 5 In chemical reactions, atoms are rearranged, separated, or combined, but are never created nor destroyed.

  9. Types of Molecules • Diatomic Molecule: contains two atoms • Homoatomic molecule: contain only one kind of atom. • Heteratomic molecule: contain more then one kind of atom • Triatomic molecule: 3 non-identical atoms • Polyatomic molecule: more then 3

  10. Classification of Matter • Pure Substances - not adulterated or mixed with anything else. • They have unique and consistent physical and chemical properties. • Physical Properties: Melting Point Temperature, Color, Density. • Chemical Properties: a chemicals ability to react with other pure substances. In a chemical reaction, substances lose their identity and form new substances with new chemical and physical properties. • It undergoes physical change without losing its identity... Eg.( melting, freezing, or evaporation)

  11. Mixtures • Consists of two or more pure substances in varying proportions. • Heterogeneous – visibly discontinuous..like salt and pepper. • Homogeneous – have a uniform appearance throughout; like sugar and water. The mixture is called a solution, and it is described as homogeneous. • Mixtures can be separated back into their pure substance components. • Mixtures have properties that are variable and depend on the proportions of the components.

  12. Compounds and Elements • Compounds: Some pure substances are found to be able to be decomposed into simpler pure substances. • Elements: pure substances that cannot be further decomposed. It cannot be separated chemically in to simpler substances, nor be created by combining simpler substances.

  13. Measurement and the Metric System • Measure: - the size, capacity, extent, volume or quantity of anything, especially as determined by comparison with some standard or UNIT. • Systeme International d’Unite’s.

  14. Significant Figures • Communicating Degrees of Uncertainty • 4 1 sig fig • 4.0 2 sig figs • 4.000 4 sig figs • 4.0000 5 sig figs

  15. Examples • Do not over represent the amt. of precision that you have. • Which digits are really giving me info about how precise my measurement is? • 0.00700 • If you measure the above in km, it could also be 7.00 m (the previous zeros are determining the units to use, the trailing zeros determine precision) • 0.052 …(could be re-written as 52 m) • (do not count leading zero’s before the first non-zero digit) • 370. (because they wrote a decimal, it is exactly 370) 3 sig figs. • 10.0 (go to nearest 10th) 3 sig figs • 705.001 ( zero’s are part of measurement ..between non zero digits) • 37,000 (ambiguous) Maybe you measured to the nearest 1000, or nearest 1…you don’t know. Go with 2 sig figs. (more conservative) • ( A trailing Zero as in 4.130 is significant. (This has 4 sig figs)…

  16. Rules of Thumb • A trailing zero , 4.130 , is significant. • A zero within a number, 35.06 cm • A zero before a digit as in 0.082 , is not significant • A number ending in zero with no decimal point , 20 is ambiguous.

  17. Mulitiplication and Division with Sig Figs • Let’s say we are calculating the area of a Rectangle 1.69 m x 2.09m • Area = 3.5321 m2 Use the least precisice number as the basis for the amt. of sig figs.. 3 sig figs.. • Area = 3.53 m2

  18. Another Example • Calculate how many tiles I need for a room 12.07 ft x 10.1 ft. • Floor Area = 121.907 ft2 • ** Do not round yet! *** go to the end with all numbers, then establish sig figs and round** • Tiles in bathroom = 121.907ft2/1.07 ft2 • Tiles = 113.931775701 tiles • 3 sig figs = 114 tiles

  19. Addition and Subtraction • Ex: 1.26 (nearest hundredth , 3 sig figs) + 2.3 (nearest 10th, 2 sig figs) = 3.56 • The least precise number went to the 10th, therefore 2 sig figs in result. 3.7 • Or 1.26 + 102.3 = 103.56 (only as precise as the least precise number)… • 1.26 has 3 sig figs, 102.3 has 4 sig figs, however the least precise measurement is 102.3 as it is measured only to the 10th, not the 100th…therefore the answer will be 4 sig figs, 103.6

  20. Another Example • One Block: 2.09 m high • Another block: 1.901 m high • How tall is it to stack them? • 1.901 + 2.09 = 3.991 • Did I measure the entire stack to the nearst mm? NO! Only report as precise as the least precise measurement. 3.99m

  21. Another example • Building: 350 ft tall (ambiguous) • Radio Tower : 8 ft tall • How tall is building plus the tower: 358 ft • We only measured tower to nearest ft. • You have to round to the nearest 10 ft. Answer is actually 360 ft. or report it as 3.6 x 10 2

  22. Metric System Units

  23. Using Units in Calculationsaka: Dimensional Analysis • Step 1: Write Down the known or given quantity. Include both numerical value and units of the quantity. • Step 2: Leave some working space and set the known quantity equal to the units of the unknown quantity. • Step 3: Multiply the known by one or more factors (conversion factors) to cancel the units of the known and generate the units of the unknown. • Step 4: Do the arithmetic .

  24. Example Problems • 50 μL (50 – microliter) sample of blood serum must be expressed as Liters. • (1μL = 1 x 10-6 L) • Step 1: • Step 2: • Step 3: • Step 4:

  25. Example • One of the fastest-moving impulses in the body travels at a speed of 400 ft/per second. (ft/s). What is the speed in miles/hr?

  26. Calculating Percentages • Percent = number of specific items x 100 total items in the group % = Part/Total x 100

  27. Non SI units most common

  28. Prefixes for Metric Units

  29. Mass, Volume, & Density • Mass is the measure of a quantity of matter. It is measured relative to a standard mass (which is why the devices to weigh an object are called Balances) Mass is not Weight. • Volume is the amount of space a sample occupies. 1 mL = 1cm3 • Denisty – a physical property of a substance = mass/volume. Since volume increases with Temperature increase, a density is always reported with a Temperature.

  30. Example • A 35.66 g sample of metal as weighed and put into a graduated cylinder that contained 21.2 mL of water. The water level after the metal was added was 25.2 mL. What is the density of the metal in g/cm3

  31. How to measure Density • Take a substance and weight it. Then add it to a known volume of water in a volumetric flask and notice the volume change in the water as the volume of the substance. Calculate the density. • To calculate the mass of a liquid, you add the liquid to a zeroed balance with a volumetric flask and weight the liquid. Then you have the mass and the volume, and you can caluclate the density.

  32. Density Example Problem • What is the volume of a 32 g sample of ethanol whose density is 0.789g/cm3 ? Report volume in cm3

  33. Hydrometer • If something floats on water, it is less dense then water. • If something sinks, it is more dense then water. • A hydrometer rises or falls to a density that is equal to the density of the liquid. It is calibrated to show the specific gravity of the liquid.

  34. Specific Gravity • S.G. = density of test liquid/ density of reference liquid • Note that the units cancel. • The standard reference liquid for measuring the specific gravity of aqueous solutions is pure water at 4 deg. C. Density is 1.000g/cm3 . • S.G. of blood is 1.028. • This means that blood is 1.028 times the density of pure water.

  35. Temperature • Substances can either gain heat or lose it, depending on whether they are cooler or hotter then their environments. • To measure heat, we must have an indication of how hot or cold something is…that is the temperature. • It indicates how hot something is…not the amount of heat.

  36. Kelvin, C, and F • K = C + 273 • F = 9/5 C + 32 • C (F-32)(5/9)

  37. Heat and Calorimetry • Heat is a form of energy. • Each substance has a different capacity to absorb heat. • A unit of heat is defined by its effect (the rise in temperature) on a fixed mass of a reference substance. • SI unit = joule, (J) • Non-SI commonly used = cal • 4.184 J = 1 cal

  38. Specific Heat : Cp • The characteristic response of a given mass of a given substance to a given amount of heat is expressed by Cp • Cp= joules/(grams x Δ°C) • The specific heat is equal to the heat absorbed or lost per Celsius degree change in temperature per gram of substance.

  39. Specific Heat • The higher a substances specific heat, the more slowly it’s temperature rises in repsonse to heating.

  40. Calculating Specific Heat • What is the specific heat of a substance if the addition of 334 J of heat to 52 g of that substance causes the temperature to rise from 16 C to 48 C?

  41. Calculating Heat from Cp • How much heat must be added to 45.0 g of a substance that has a specific heat of 0.151 J/gC to cause it’s t to rise from 21 C to 47 C ?

  42. Calorimeter • When the heat produced by some physical or chemical process is abosrbed into a given mass of water, the water’s T rise will allow us to calculate the heat produced by the process.

  43. Example • What is the Specific Heat (Cp) of an element that takes 50 Joules to heat up 400 grams from 34 deg to 76 deg.?

  44. Calorimeter

  45. Basal Metabolic rate • BMR is the minimum metabolic activity of a human at rest and with an empty gi tract. • In nutrition and metabolism, heat is more commonly given as calories. • The rate means it is the amount of heat over a period of time…expressed in kcal/min.

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