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Chem 111, Week 1 Notes

by: je

Chem 111, Week 1 Notes Chem 111-007

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Principles of Chemistry I
Corina Brown
Class Notes
matter, measurement, properties, chemical, Physical, mixture, pure, substance, conversions
25 ?




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This 8 page Class Notes was uploaded by je on Saturday September 10, 2016. The Class Notes belongs to Chem 111-007 at University of Northern Colorado taught by Corina Brown in Fall 2016. Since its upload, it has received 3 views. For similar materials see Principles of Chemistry I in Chemistry at University of Northern Colorado.


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Date Created: 09/10/16
August 22, 2016 Chapter 1: Introduction Matter & Measurement  Chemistry is the study of matter and its changes. o Macroscopic (observable) level o Sub-microscopic (particulate) level  Matter – anything with mass and volume  Mass – amount of substance in an object  Weight – measures the gravitational pull  Volume – amount of space taken by an object  Density – mass (g) divided by volume (cm ) m=18.96g d=18.96g/4.31cm^3 3 3 V=4.31cm d=4.39g/cm August 24, 2016 Classification of Matter (composition) Matter Pure Substance Mixture of Pure Substance Element Compound Homogeneous Heterogeneous Pure Substance  Constant composition  Element – made of the same kind of atoms  Molecules – 2 or more elements combined Compounds – 2 or more elements chemically combined Mixture – physical combination of two or more (composition varies)  Homogeneous – uniform composition o Ex: saline solution  Heterogeneous – non-uniform composition o Ex: salt and paper Seawater Homogenous mixture Helium Gas Element Sodium Chloride Compound Bottle of Soft Drink Homogenous Mixture Milkshake Homogeneous Mixture Air in a Bottle Homogeneous Mixture Concrete Heterogeneous Mixture Classification of Matter (state)  Solid – has a definite shape and volume  Liquid – has a definite volume, but has no definite shape  Gas – has no definite shape or volume Properties & Changes of Matter Properties  Physical Changes – can be observed without changing the identity of the substance o Ex: density, boiling point, mass, volume  Chemical Changes – describes the ability of a substance to combine with or change into one or more substances o Ex: flammability, corrosiveness, reactivity with acid, gas formation Ironing Physical Rusting Nails Chemical Fireworks Chemical Frying an Egg Chemical Ice Melting Physical Bread with Butter Physical Types of Changes  Physical – do not change the composition of a substance o Ex: changes in state, temperature, volume, etc.  Chemical – changes result in new substance o Ex: combustion, oxidation, decomposition, etc. Types of Properties  Intensive – does not depend on amount of substance present o Ex: density, boiling point, color, etc. 2  Extensive – does depend on the amount of substance present o Ex: mass, volume, energy Physical/Chemical Properties 1 Ice water has a temperature of 0°C, water boils at 100°C. Physical. 2 Silver tarnish due to its ability to combine with sulfur. Chemical. 3 Nitroglycerine explodes. Chemical. 4 Acids are sour; bases are bitter. Physical. 5 Copper conducts electricity; diamond doesn’t. Physical. 6 Neon because it does not react to anything. Chemical. Separation of Mixtures (Lab information)  Physical means it can be used to separate a mixture into its pure components. o Magnet o Filtration – solid substances are separated from liquids and solutions.  Distillation uses differences in the boiling points of substances to separate a homogeneous mixture into its components.  Chromatography – this technique separates substances on the basis of differnces in solubility in a solvent. Mass kg 1kg=1000g length m 1km=1000m 1g=1000mg 1m=100cm 1mg=1000µg 1m=1000mm 1kg=2.2lbs 3 L x L = Area L x L x L = Volume Convert 1.5km to m Convert 15dL to L 1.5km x 1000m = 1500m 15dL x 1 L = 1.5L 1km 10dL 4 August 26, 2016 Application – Displacement Method m=12.0g d = m/v V initial = 10mL V = 16mL-10mL d = 12g/6mL V final = 16mL V = 6mL d = 2g/mL Heat vs. Temperature  Heat : energy that is transferred from hotter objects to cooler objects  Temperature : the average kinetic energy of a sample o Kelvin is the SI temperature scale o Celsius and Kelvin scales are most often used o Fahrenheit scale is not used in scientific measurements 0°C = 273K = 32°F K = °C + 273.15 °F = (9/5)(°C)+32 °C = (5/9)(°F – 32) 1. 25°C to K K = 25°C + 273.15 K = 298.15 2. 273K to °C 273K = °C + 273.15K °C = 0 or -.15 3. 179.2°F to °C °C = (5/9)(179.2 – 32) °C = 81.78°C Uncertainty in Measurements  The uncertainty in measurement is determined by the measuring device Measurements  Every measurement carries a degree of uncertainty or error  Exact Number o Has a number that is known exactly  Inexact numbers o Depend on how they are determined o Scientific instruments have limitations Accuracy vs. Precision  Accuracy o Refers to the closeness of a measurement to the true value of the quantity  Precision o Refers to the closeness of several measurements to each other Significant Figures (Sig Figs)  Sig Figs are the number of digits that are known to be accurate plus one more  The number of sig figs is determined by the measuring device Sig Fig Rules 1. All digits from 1 through 9 are significant 2. Zeroes between two sig figs are significant 3. Zeroes at the beginning of a number are never significant 4. Zeros at the end of a number are significant if a decimal point is written after the number 5. When a number ends with zeros but contains no decimal point, zeroes are not significant 2 There are practice problems on the worksheet handed out in class today. 3


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