Chm 113 Quiz 1
Mixtures
(1 or more substances) Homogeneous (uniform throughout) Heterogeneous (not uniform throughout) two or more pure substances that exist together but are not combined chemically. can be separated by physical processes Important: Mixtures can have properties distinct from any of the individual substances making up the mixture.
Ernest Rutherford
(1910) *The nuclear Model of the Atom*: He shot alpha particles at a thin sheet of gold foil and observed a unique pattern of scatter of the particles. Rutherford postulated a model of the atom with a very small, positively-charged, dense nucleus with the electrons around the outside. *Rutherford's theory lead to his discovery of protons * in 1919. Neutrons were discovered by James Chadwick in 1932. Because the nucleus is small, dense and positively charged the alpha particles
Physical Changes
(Physical Properties = physical changes) These are changes in matter that do not change the composition of a substance. Changes of state, temperature, volume, etc. (Example: H2O(l) --> H2O(g)) (Ex: Changes of state (phase changes) this includes freezing, melting, sublimation (dry ice melts into a gas at 25 C), and deposition)
Group 1A
*Alkali metals; Elements: Li, Na, K, Rb, Cs, Fr; soft reactive metal* (does not include H)
Group 2A
*Alkaline Earth metals; Elements: Be, Mg, Ca, Sr, Ba, Ra*
Group 6A
*Chalcogens; Elements: O, S, Se, Te, Po*
Group 7A
*Halogens; Elements: F, Cl, Br, I, At*
Group 8A
*Noble Gases (Rare gases); Elements: He, Ne, Ar, Kr, Xe, Rn. Tend to be a nonreactive gas*
Ions
*atoms* or *groups of atoms* that have an electrical charge, called *cations* and *anions.* Certain elements have a predictable charge based on their position in the periodic table. *(Zn2+, Ag+, Cd2+)*
Metalloids
*border the stair-step line (with the exception of Al, Po, and At). Their properties include: Electronegativities between those of metals and nonmetals, Ionization energies between those of metals and nonmetals, Possess some characteristics of metals/some of nonmetals, Reactivity depends on properties of other elements in reaction, and Often make good semiconductors*
Metals
*on the left side of the chart. They include main group metals and transition metals. Their properties include: Shiny 'metallic' appearance, Solids at room temperature (except mercury), High melting points, High densities, Large atomic radii, Low ionization energies, Low electronegativities, *Lose electrons easily*, Malleable Ductile, Thermal conductors, Electrical conductors*
Nonmetals
*on the right side of the periodic table (with the exception of H). Their properties include: High ionization energies, High electronegativities, *Gain electrons easily*, Poor thermal conductors, Poor electrical conductors, Brittle solids, and Little or no metallic luster*
Mass of a Single Electron
*𝐸𝑙𝑒𝑐𝑡𝑟𝑜𝑛 𝑀𝑎𝑠𝑠 = (1.602 × 10^−19𝐶)/(1.76 × 10^8𝐶/𝑔) = 9.10939 x 10^-28 g* Known today to 6 sig figs. Mass is about 2000 times smaller than that of hydrogen, the lightest atom. If you had a bag of electrons weighing 1 lb, how many electrons would be in the bag? 4.98 x 10^29 electrons
Pure Substances
1 substance: Elements, Compounds, Molecules. have definite compositions, and each has a set of properties that are unique.
The Discovery of Atomic Structure Cathode Rays & Electrons
1897—*J.J.Thomson discovered the electron* using a cathode-ray tube by detecting a negatively charged particle. He figured out that the particles were negative by using a magnet (the positive side attracts; the negative side retracts) (don't memorize years or how the cathode ray works only learn who discovered it, what was important, and the experiment name) *Charge:mass of an electron: 1.76 X 10^8 C/g* 1909: *Millikan* used the *oil-drop experiment* to calculate the charge of an electron. *He discorered that the charge of the electron was 1.6 X 10^-19 C*
Dalton's Theory
4 postulates: (1) Each element is composed of extremely small particles called atoms. (2) All atoms of a given element are identical, but the atoms of one element are different from the atoms of all other elements. (3) Atoms of one element cannot be changed into atoms of a different element by chemical reactions; atoms are neither created nor destroyed in chemical reactions. (this disproved the ideas of Alchemy of making something into gold. Sorry Rumpelstiltskin!) (4) Compounds are formed when atoms of more than one element combine; a given compound always has the same relative number and kind of atoms *Dalton's theory explains several laws of chemical composition that were known at the time. These included: Law of Conservation of Mass (postulate 3), and Law of Multiple Proportions (deduced from four postulates)*
Conversion factors
A ratio: often exact numbers and are not measurements (infinite sig figs: NEVER limit your significant figures based on conversion factors)
chromatography
Ability to adhere to a surface or difference in solubility in a solvent (EX: Pen ink lab)
distillation
Ability to form gases (differences in boiling point) (EX: separate salt (NaCl) and water (H2O)) (Fancy way of boiling)
significant figures rules
All nonzero digits are significant. Ex: 134 Zeroes between two significant figures are themselves significant. Ex: 2005 Zeroes at the beginning of a number are never significant. Ex: 005 or 0.07 Zeroes at the end of a number are significant if a decimal point is written in the number. Ex: 2.200 (4 sig figs) When reading a measurement number, to determine the number of significant figures, read the number from left-to-right starting with the first non-zero digit.
The Scientific Method
Allows us to think critically: Ask a question, do background research, construct a hypothesis, test an experiment, analyze results and draw conclusion, if the hypothesis is True report the results, if they are False or partially true think about it again and construct a new hypothesis
Inexact Numbers
Any number obtained by measurement, uncertainties *always* exist in measured quantities (will always include uncertainties)
Intensive Properties
Are Independent of the amount of the substance that is present (useful for identifying substances). (Ex: Density, boiling point, color, etc.) It doesn't matter how much there is it will still remain the same
Extensive Properties
Are dependent upon the amount of the substance present. (EX: Mass, volume, energy, etc.)
Physical Properties
Can be observed without changing a substance into another substance. (Ex: Melting point of aluminum,Density of mercury, Boiling point, density, mass volume, shiny, etc.)
Chemical Properties
Can only be observed when a substance is changed into another substance. (flammability, Ability of nitric acid to dissolve copper, etc.)
Chemical Changes
Chemical changes result in new substances. (Combustion, oxidation, decomposition, etc.) Example: 2H2O (g) --> 2H2 (g) + O2 (g)). (water is decomposed into Hydrogen and Oxygen gas. Ex: Converting elements to compounds Combustion of hydrogen: In the course of a chemical change, the reacting substances are converted to new substances. Evidence: Gas generated Heat evolved Color change
The Atomic & Molecular Perspective
Chemistry explains properties & behavior of matter based on structure & events at the *atomic and molecular levels* (the microscopic levels)
Separation of Mixtures
Components of a mixture can be separated by *physical* processes. These processes take advantage of the differences in physical properties of the components.
Homogeneous Mixtures (also called solutions)
Constant composition throughout. Solutions are often liquids but they can be mixtures of solids and gases as well. You cannot pick out the different parts (Ex: Closed soda the gas and the liquid are constant), air, salt water(aq), metal alloys, sugared tea)
Liquid
Disorder, particles or clusters of particles are free to move relative to each other; particles are closer together. When they cool down they will become a solid
Atomic Number
Each specific element has a fixed number of protons. The number of protons. Atomic Number (Z) = the number of protons or electrons (EX: Carbon = 6 protons).
Exact Numbers
Exact counts of objects: Values which are known exactly, or defined values. (Ex: 12 eggs = 1 dozen eggs, 2.54 cm = 1 inch, 1 m = 100 cm)
Law of Multiple Proportions:
If two elements A and B combine to form more than one compound, the masses of B that can combine with a given mass of A are in the ratio of small whole numbers. (EX: For H2O = 16g O for every 2 g H for a ratio to 8 to 1 (see periodic table) For H2O2 = 32 g O for every 2 g H for a ratio to 16 to 1) From this we can conclude that H2O2 contains twice as many atoms of O per H than H20
Law of Constant Composition
In a given compound, the relative numbers and kinds of atoms are constant. (EX: H2O: 1 part H, 8 parts O, by mass-ALWAYS!) Elements combine in specific proportions (EX: According to the Law of Constant Composition what is the ratio of CO2: 1 part C, 2.7 parts O, by mass - ALWAYS)
Mass Number
In a neurtral atom the number of electrons equals the number of protons. The number of neutrons can vary Number of protons plus neutrons
The Discovery of Atomic Structure The Nuclear Model of the Atom
J.J. Thomson: proposed the "plum pudding" model of the atom. It featured a positive sphere of matter with negative electrons imbedded in it. This model was later disproved by Rutherford
Early Atomic Theory
Lots of evidence was needed to show that matter is composed of the atoms we cannot see. The theory that atoms are the fundamental building blocks of matter reemerged in 1808 John Dalton published a paper on atomic theory. The first person to come up with the atom is Democratis but Dalton had backing for his ideas
Magnetic Properties
Mixture of iron and sulfur.
Crystalline Solid
Ordered arrangement; particles are essentially fixed in positions; particles are close together.
Protons, Neutrons and Electrons
Proton (+ charge) with a mass of 1.0073 amu. Neutron (no charge) with a mass of 1.0087 amu. Electron (- charge) with a mass of 5.486 X 10^-4 (Learn the charge not the mass) Protons (+) and electrons (−) are the only atomic particles that have a charge. The charge of a proton is opposite in charge but equal in magnitude to that of an electron. Neutrons are neutral (hence, the name). (they are like the glue in the nucleus) Protons and neutrons have essentially the same mass. Electrons are much smaller. Every atom has an equal number of electrons and protons, so atoms have no net electrical charge.
Changes of state
Solid<--> liquid (Freezing <--> melting) liquid <--> gas (condensing <--> boiling (evaporation) solid <-->gas (Deposition <--> sublimation)
Properties of Matter
Substances can exist in the solid, liquid, or gas state.
SI Units
Système International d'Unités (particular choice of metric units for use in scientific measurements) A different *base unit* is used for each quantity.
Law of Constant Composition (Definite Proportions)
The elemental composition of a compound is always the same. (Ex: Pure water from any source is composed of two H atoms and one O atom.) (NH3, CH4)
Adding and Subtracting
The final answer is expressed with the same number of decimal places as the measurement with the fewest decimal places. (EX: 89.332 +1.1 = 90.4)
Multiplying and Dividing
The final answer is expressed with the same number of significant figures as the measurement with the *fewest* significant figures. (4.4643 X 2.8 = 13)
Formula Unit
The ionic compound "empirical formula"
Periodic Table
The periodic classification of elements was developed in 1869 by Dmitri Mendeleev. Elements are arranged in order of atomic number. The rows on the periodic chart are periods. Columns are groups (or families). - Elements in the same group have similar physical & chemical properties. When you look at the chemical properties of elements, you notice a repeating pattern of reactivities. The elements can also be classified as Metals, Nonmetals, and Metalloids
Radioactivity
The spontaneous emission from certain elements of high-energy particles and radiation (1896 Becquerel & Curie). Three types of radiation were later discovered by Ernest Rutherford (1896-1908): alpha particles (positive charge), beta particles (negative charge), and gamma rays (no charge)
The Atomic Theory of Matter
There are over *100* distinct elements. Each element is composed of unique atoms. Think of each type of atom as one of the letters in the English alphabet. There are only 26 letters in the alphabet but there are over 700,000 words in the English language. How many different chemical compounds do you think could be formed from 100 elements? A lot
Gas
Total disorder; much empty space; particles have complete freedom of motion; particles far apart. If the gas are in a much smaller volume they will interact more and get more excited which means they can stick together. If they cool down they will become a liquid
Empirical Formulas from Analysis
We can determine the empirical formula for a compound from the % composition: % Composition (mass ratio)→Empirical formula (mole ratio). In order to do this 1. Assume 100 g of compound. 2. Convert grams of each element to moles. 3. Divide or multiply each by the same number to get whole-number subscripts. (EX: Given that there is 30.45% N and 69.56% O what is the formula for the compound? 30.45 g N X 1 mol N/14.01 = 2.173 (limiting reactant) 69. 56 O X 1 mol O/ 16 = 4.348 2.173/2.173 = 1 N, 4.348/2.173 = 2 O There are 2 O per 1 N so it is NO2)
The Modern View of Atomic Structure
While protons and neutrons make up most of the mass of an atom, the electrons make up the volume. (the sheer volume held by the negative atoms are enormous compared to the nucleus)
Pure substance
a homogenous compotision that can either be a compound (H20) or an element (Ag)
temperature
a measure of the average kinetic energy of the particles in a sample. the Celsius and Kelvin scales are most often used. If something is cold there is not a lot of movement (g) but is something is hot there is a lot of movement (g)
Density
a physical property of a substance. (INTENSIVE Property) It has units that are derived from the units for mass and volume (EX: g/mL, g/cm3) d = m/V (capital V = volume)
Matter
anything that has mass and occupies space
Isotopes
are atoms of the same element with different masses. Isotopes have different numbers of *neutrons*. Never change the amount of protons. Often times only the mass number and charge is written for a specific isotope (EX: 14C: Protons & Electrons = 6; number of neutrons = 8)
covalent bonds
bonding through sharing of electrons between atoms. The subscript (subscripts indicate how many are in each molecule and are like body parts, it would be hard to add another arm or take one off) to the right of each element symbol indicates the number of atoms of that element in one molecule of the compound. *The order in which a molecular formula is written does not indicate the order of the bonded atoms.*
Compounds*
combinations of atoms of more than one type of element.
Molecular compounds
composed of atoms chemically bonded together. The elements in molecular compounds are almost always *nonmetals*
Diatomic elements
composed of molecules that contain two identical atoms. Know these 7 diatomics: Hydrogen, oxygen, nitrogen, fluorine, chlorine, Bromine, and Iodine
Molecular elements
composed of only one type of atom.
Atomic mass unit (amu)
exactly one-twelfth the mass of a carbon-12 atom. 1/12 mass of 12C atom = 1 amu. Mass of one 12C atom = 12 amu (Exactly. This is the standard.) 1 amu = 1.66054 x 10-24 g What about the fact that different isotopes of an element have different numbers of neutrons, and hence, have different weights? Atomic mass on periodic table is the "weighted" average of abundance of an atom's isotopes. For naturally occurring carbon: (12 amu * 0.9889) + (13.0034 amu * 0.0111) = 12.011 amu
Molecular formulas
give the exact number of atoms of each element in a compound.
Empirical formulas
give the lowest *whole- number* ratio of atoms of each element in a compound. (What is the empircal formual for C6H6 = CH)
Molecules*
group of atoms (two or more) bonded chemically. (not all molecules are compounds: H2, O2) they are the same element
Anions
have a negative charge. Gain of electrons results in anions: Non-Metals typically *gain* electrons to form anions.
Cations
have a positive charge. Loss of electrons results in cations: *Metals typically lose electrons to form cations*
Solubility
iron dissolves in acid but gold does not.
Volume
length X width X height (1 cm^3 = 1 mL) We use a variety of containers to measure & dispense liquids in chemistry.
element
made of the same kind of atom.
molecule
made of two or more atoms joined together (oxygen)
compound
made of two or more different kinds of elements. (CO2 ethylene glycol, and sodium chloride)
Law of Conservation of Mass
mass is neither created nor destroyed in a chemical reaction. The total mass of the materials present after a chemical reaction is the same as the total mass present before the reaction
Heterogeneous Mixtures
non-uniform composition you can easily separate the two (Ex: opened Soda (the gas and the liquid are no longer constant) , potting soil, oil and water, granite, iron and sulfur)
pure substances
not mixed with other compounds or elements (Examples: Sodium Chloride (NaCl), Ammonia (NH3), Water (H2O), Iron(III) oxide (Fe2O3) (rust), Iron(III) sulfide (Fe2S3))
Fahrenheit
not used in scientific measurements. *The freezing point is 32 F, boiling point is 212 F and normal body temperature is 98.6 F (180 degree intercal Interconvert C & F: F = 9/5( C) + 32, C = 5/9( F − 32)
Metal cation
polyatomic anion compounds are also held together by electrostatic attraction between the + and − charges.
significant figures
refers to digits that were measured. The number of significant digits describes the exactness of the measurement. A measured quantity expressed as a number will have digits that are certain plus one uncertain digit (the right-most digit will always have uncertainties). When rounding calculated numbers, we pay attention to significant figures so we do not overstate the accuracy of our answers (We want to estimate to three significant digit.) *The last digit in your measurement should have some uncertainty*
Accuracy
refers to the proximity of a measurement to the true value of a quantity. (Similar to accepted value but not necessarily close together)
Precision
refers to the proximity of several measurements to each other. (Similar to each other but not necessarily close to the accepted value)
Celsius
scale is based on the properties of water. *0 C is the freezing point of water. 100 C is the boiling point of water and 25 C is room temperature Lowest possible temperature -273.14 C*
Structural formulas
show the order in which atoms are bonded (the connectivity)
Kelvin
the SI unit of temperature and is based on the properties of gases. *There are no negative in Kelvin temperatures! Absolute zero: The lowest possible T in K is 0 K boiling point is 373 K, freezing point is 273 K and 298.15 K is normal body temperature* (-273.15 C) K = C + 273.15
Atoms
the building blocks of matter.
Diatomic molecule
two atoms in the molecule (oxygen)
Phase
use filtration to separate a solid phase from a liquid phase.
Measurements
used to communicate quantitative information. A measurement consists of (1) a number (2) an appropriate unit. (Ex: 125 C; 10.0 kg; 10.0 mg; 1.5 yards; 1.004 meters; 0.936 g/mL; 10 cm3)
Dimensional Analysis**
used to convert one quantity to another. used when solving problems in which some unit(s) of measure, e.g., meters, grams, is converted into other units of measure.
Exponential (Scientific) Notation
used to express small & large numbers conveniently. Example: SI base unit of length is meter but typically we deal with much smaller units of length in chemistry. The radius of a sodium atom is 1.54 x 10−10 m = 154 pm Example: SI base unit of mass is kilogram but typically we deal with much smaller units of mass. The mass of a molecule of glucose is 2.99 x 10−25 kg = 2.99 x 10−22 g
