Chemistry Chapters 1-2

Unit of measurement

A measurement (a.k.a. quantitative observation) is only meaningful when it consists of both a NUMBER and a SCALE

Density

Density= mass/volume

Homogenous Mixture (Solution)

Has visibly indistinguishable parts; composition the same throughout. [Air (a mixture of gases); Wine (a complex liquid solution); Brass (a solid solution of copper and zinc)]

Matter: Anything occupying space and having mass.

Mass: measure of the quantity of matter. (kg) Weight: force that gravity exerts on an object. (W=m x g; W=m x c) 3 states: Solid, liquid, and gas

Mixtures

Most of the matter around us consists of pure substances.

Solid

Rigid; Fixed volume and shape

Chromatography

The general name applied to a series of methods that employ a system with two phases (states) of matter; a mobile phase (liquid or gas) and a stationary phase (solid)

Diatomic molecule

^^ Two chemical elements

Uncertainty in Measurement

a measurement always has some degree of uncertainty. The certain digits together with the uncertain digit are the significant figures of a measurement

Macroscopic world

we experience

Molecular Compounds

•Contain discrete molecular units. •Usually composed of nonmetallic elements; Nonmetal + Nonmetal •Many are binary compounds. •Naming binary molecular (or binary covalent) compounds (Type III): is similar to naming binary ionic compounds. •Name the first element in the formula first; element furthest to the left in a period and closest to the bottom of a group on periodic table is placed first in formula •Second element named by adding "-ide" to the root of the element name •Examples: •HCl - Hydrogen chloride •HBr - Hydrogen bromide •SiC - Silicon carbide

Subatomic Particles

•Electrons •Protons •Neutrons

Ionic Compounds

•Many ionic compounds are binary compounds, compounds formed from just two elements. •Naming a binary ionic compound (Type I): •First Element Named: metal cation •Second Element Named: nonmetal anion; add "-ide" to element name.

The Electron

•Much information was learned about atomic structure through the study of radiation - the emission and transmission of energy through space in the form of waves. •J. J. Thompson used a cathode ray tube to determine the ratio of electric charge to the mass of an individual electron is -1.76 108 C/g [C stands for coulomb, which is the unit of charge.]

The Neutron

•Rutherford's model did not account for all the subatomic particles. •This was revealed when comparisons were made of the ratio of the mass of a helium atom to that of a hydrogen atom, which should have been 2:1 based on Rutherford's model. •However, it was found to be 4:1, indicating that there must be one or more other particles in the nucleus not accounted for. •Chadwick discovered neutrons, which are electrically neutral particles having a mass slightly greater than that of protons. •Hence, the mass ratio mystery was explained: •The nucleus of helium has 2 protons and 2 neutrons, whereas, the nucleus of hydrogen only has one proton and zero neutrons; therefore, the ratio is 4:1.

The Proton and the Nucleus

•Two clear features about atoms known by the early 1900s: 1.Atoms contain electrons 2.Atoms are electrically neutral •In order to be electrically neutral, then there must be an equal number of positive and negative charges. •Thompson's "Plum Pudding" Model: an atom is a uniform, positive sphere of matter in which electrons are embedded.

Atom

The smallest part of an element that is still that element

Molecule

Two or more atoms joined and acting as a unit

A & B

Two or more oxoacids can have the same central atom but a different number of O atoms. •Using the reference oxoacids, the following rules are used to name these compounds. 1.Addition of one O atom to the "-ic" acid: The acid is called "per...-ic" acid. [Example: adding an O atom to HClO3 changes chloric acid to perchloric acid, HClO4] 2.Removal of one O atom from the "-ic" acid: The acid is called "-ous" acid. [Example: removing an O atom from HNO3 changes nitric acid to nitrous acid, HNO2] 3.Removal of two O atoms from the "-ic" acid: The acid is called "hypo...-ous" acid. [Example: removing two O atoms from HBrO3 changes bromic acid to hypobromous acid, HBrO]

Two Common Isotopes of Uranium

Uranium-235, Uranium-238 With the exception of hydrogen, isotopes of elements are identified by their mass numbers. •The chemical properties of an element are determined primarily by the protons and electrons in its atoms. •Neutrons do not take part in chemical changes under normal conditions. •Hence, the isotopes of the same element have similar chemistries, forming the same types of compounds and displaying similar reactivities.

Chemical elements

Water decomposes to hydrogen and oxygen

microscopic world

atoms and molecules

Distillation

one of the most important methods for separating the components of a mixture; the process depends on differences in the VOLATILITY of the components.

SI unit

picometer 1pm=1x10^-12m

MM prefix

tera- T 10^12 1 terameter (™) = 1 × 10^12 m giga-G 10^9 1 gigameter (Gm) = 1 × 10^9 m mega-M 10^6 1 megameter (Mm) = 1 × 10^6 m kilo- k 10^3b1 kilometer (km) = 1 × 10^3 m deci-d 1/10 or 10^−1b1 decimeter (dm) = 0.1 m centi-c 1/100, or 10^-2 1 centimeter (cm) = 0.01 m milli-m 1/1,000, or 10^−3 1 millimeter (mm) = 0.001 m micro-μ 1/1,000,000, or 10^−6 1 micrometer (μm) = 1 × 10−6 m nano-n 1/1,000,000,000, or 10^−9 1 nanometer (nm) = 1 × 10^−9 m pico-p 1/1,000,000,000,000, or 10^−12 1 picometer (pm) = 1 × 10^−12 m

Law of conservation of Mass

the observation that the total mass of materials is not affected by a chemical change in those materials

Law of Multiple Proportions

•Another law supported by the third hypothesis is the Law of Multiple Proportions - if two elements can combine to form more than one compound, the masses of one element that combine with a fixed mass of the other element are in ratios of small whole numbers. •As explained by Dalton, the compounds differ in the number of atoms of each kind that combine.

•Law of Multiple Proportions Example

•Carbon can form 2 stable compounds with oxygen - carbon monoxide & carbon dioxide. •Carbon monoxide = 1 atom of carbon + 1 atom of oxygen. •Carbon dioxide = 1 atom of carbon + 2 atoms of oxygen. •Ratio of oxygen in carbon monoxide to oxygen in carbon dioxide is 1:2. •The result is consistent with the Law of Multiple Proportions, because the mass of an element in a compound is proportional to the number of atoms of the element present. if two elements can combine to form more than one compound, the masses of one element that combine with a fixed mass of the other element are in ratios of small whole numbers.

Acids and Bases

•Example: HCl - Hydrogen chloride or hydrochloric acid...depending on physical state. •Gaseous or pure liquid: it's a molecular compound called hydrogen chloride •When dissolved in water:

Formulas of Ionic Compounds

•Ionic compounds consist of a combination of cations and anions. •The formula is usually the same as the empirical formula. •The sum of the charges on the cation(s) and anion(s) in each formula unit must equal zero.

Two fundamental concepts of Chemistry

1. Matter is composed of various types of atoms 2. One substance changes to another by reorganizing the way atoms are attached to each other

Dalton's Atomic theory

1.Elements are composed of extremely small particles called atoms. 2.All atoms of a given element are identical, having the same size, mass and chemical properties. The atoms of one element are different from the atoms of all other elements. 3.Compounds are composed of atoms of more than one element. In any compound, the ratio of the numbers of atoms of any two of the elements present is either an integer or a simple fraction. 4.A chemical reaction involves only the separation, combination, or rearrangement of atoms; it does not result in their creation or destruction.

Three Categories of Elements

1.Metals: a good conductor of heat and electricity 2.Nonmetal: a poor conductor of heat and electricity 3.Metalloid: has properties that are intermediate between those of metals and nonmetals. •The majority of known elements are metals. •17 classified as nonmetals •8 classified as metalloids

Conversion(liter)

1L=1000mL =1000cm^3 =1dm^3

Conversion(mass)

1kg=1000g=1 x 10^3g

Conversion(liter)

1mL=1cm^3

bases

A base can be defined as a substance that yields hydroxide ions (OH-) when dissolved in water. •Ammonia (NH3) is a molecular compound in the gaseous or pure liquid state and is a common base. •When ammonia dissolves in water, NH3 reacts partially with water to yield 𝐍𝐇𝟒+ and 𝐎𝐇− ions.

Physical Change

A change in the FORM of a substance, not in its chemical composition; does not alter the composition or identity of a substance. Can be used to separate a mixture into pure compounds. Will not break compounds into elements.

Electrolysis of Water

A chemical process that uses an electric current passed through water to break down the water molecules into free elements hydrogen and oxygen (chemical change)

Natural Law

A concise statement of a relationship between phenomena that is always the same under the same conditions. Summarizes WHAT happens.

Molecules

A diatomic molecule contains only two atoms: •Hydrogen (H2) •Nitrogen (N2) •Oxygen (O2) •Bromine (Br2) •Iodine (I2) •Hydrogen chloride (HCl) •Carbon monoxide (CO) •A polyatomic molecule contains more than two atoms: •Ozone (O3) •Water (H2O) •Ammonia (NH3) •Methane (CH4)

Empirical Formulas

A molecular formula shows the exact number of atoms of each element in the smallest unit of a substance; this is the true formula of a molecule. •Example: Hydrogen peroxide - H2O2 •Each hydrogen peroxide molecule consists of 2 hydrogen atoms and 2 oxygen atoms. •Ratio of hydrogen to oxygen atoms is 2:2 or 1:1. •An empirical formula shows the simplest whole-number ratio of the atoms in a substance; written by reducing the subscripts in the molecular formulas to the smallest possible whole numbers. •For hydrogen peroxide, the empirical formula is HO. •The empirical formula can be obtained from the molecular formula, but not vice versa. •The empirical formula is usually determined first when analyzing an unknown compound; the molecular formula is deduced next. •The molecular formula and empirical formula are many times the same: water (H2O), ammonia (NH3), carbon dioxide (CO2), and methane (CH4).

Ions

A monatomic ion contains only one atom: •𝐍𝐚+ •𝐂𝐥− •𝐂𝐚𝟐+ •𝐎𝟐− •𝐀𝐥𝟑+ •𝐍𝟑− •A polyatomic ion contains more than one atom: •𝐎𝐇− •𝐂𝐍− •𝐍𝐇𝟒+ •𝐍𝐎𝟑−

Density

A property of matter representing mass per unit. Common units are g/cm^3 and g/mL

Element

A substance that CANNOT be decomposed into simpler substances by chemical or physical means.

Compound

A substance with constant composition that can be broken down into elements by s chemical processes.

Acids and Bases

Acid: a substance that yields hydrogen ions (H+) when dissolved in water. •H+ is equivalent to one proton and is often referred to way. •Formulas for acids have one or more hydrogen atoms and an anionic group. •Anions that end in "-ide" have associated acids with a "hydro-" prefix and an "-ic" ending.

Theory (Model)

An EXPLANATION of behavior. An attempt to explain WHY something happens.

Ions

An ion is an atom, or group of atoms, that has a net positive or negative charge. •In chemical reactions, the negatively charged particles (electrons) can be lost or gained. •Cation: an ion with a positive charge; if a neutral atom loses one or more electrons it becomes a cation. Anion: ion with a negative charge; if a neutral atom gains one or more electrons it becomes an anion. •Ionic compound: a compound formed from cations and ions. •Example: •Sodium chloride: 𝐍𝐚(+)+𝐂𝐥(−)→𝐍𝐚𝐂𝐥

Acids and bases

An oxoacid is an acid that contains hydrogen, oxygen, and another element. •Formula usually written with H first, then the central element, then O. •Five Common Acids Used as References in Naming Oxoacids (Note: all end with "-ic") •H2CO3 - Carbonic acid •HClO3 - Chloric acid •HNO3 - Nitric acid •H3PO4 - Phosphoric acid •H2SO4 - Sulfuric acid

Filtration

Another method of separation; used when a mixture consists of a solid and a liquid mixture is poured onto a mesh, such as filter paper, which passes the liquid and leaves the solid behind.

Atomic Number, Mass Number, and Isotopes

Atoms are identified by the number of protons and neutrons they contain. •Atomic Number (Z): the number of protons in the nucleus of each atom of an element. •Neural Atom: # protons = # electrons; therefore, the atomic number also indicates the number of electrons in the atom. •The atomic number gives the chemical identity of an atom: •Example: Nitrogen has an atomic number of 7. Therefore, neutral nitrogen has 7 protons and 7 electrons. From another perspective, any atom with 7 protons is nitrogen.

Celsius

C = (F - 32) * 5'C/9'F

Physical methods

Can be used to separate mixtures into pure substances (One with constant composition).

Chemical Formulas

Chemical Formulas: used to express the composition of molecules and ionic compounds in terms of chemical symbols.

Matter

Composed of tiny particles called atoms

Law of Conservation of Mass

Dalton's Fourth Hypothesis: A chemical reaction involves only the separation, combination, or rearrangement of atoms; it does not result in their creation or destruction. •This hypothesis gives us another way of stating the Law of Conservation of Mass - matter can be neither created nor destroyed. Because matter is made up of atoms that are unchanged in a chemical reaction, then mass must be conserved as well.

Dalton

Dalton's Third Hypothesis: Compounds are composed of atoms of more than one element. In any compound, the ratio of the numbers of atoms of any two of the elements present is either an integer or a simple fraction. •This third hypothesis suggests to us that in order to form a certain compound we need: (1) atoms of the right kinds of elements and (2) specific numbers of these atoms. •This idea is an extension of the Law of Definite Proportions - different samples of the same compound always contain its constituent elements in the same proportion by mass.

Scientific notation Division

Division 𝟖.𝟓×𝟏𝟎𝟒÷𝟓.𝟎×𝟏𝟎𝟗= 1.Divide N1 and N2. 2.Subtract exponents n1 and n2. 𝟖.𝟓÷𝟓.𝟎×𝟏𝟎𝟒−𝟗= 𝟏.𝟕×𝟏𝟎−𝟓

Fahrenheit

F = C * 9'F/5'C + 32'F

The PTable

Group 1A elements are called Alkali Metals •Group 2A elements are called Alkaline Earth Metals •Group 7A elements are called Halogens •Group 8A elements are called Noble Gases (or Rare Gases)

Liquid

Has definite volume but no specific shape. Assumes shape of container

Gas

Has no fixed volume or shape. Takes on shape and volume of its container Highly compressible.

Heterogeneous Mixture

Has visibly distinguishable parts; composition not uniform throughout. [sand in water, iced tea w/ ice cubes; cement; iron filings in sand] Can usually be separated into two or more homogenous mixtures or pure substances. [ice cubes be separated from tea]

Hydrates

Hydrates are compounds that have a specific number of water molecules attached to them

Scientific notation

If the decimal point has to be moved to the left, then n is a positive integer; if it has to be moved to the right, then n is a negative integer.

Isotopes

In most cases atoms of a given element do not all have the same mass. •Isotopes: atoms that have the same atomic number but different mass numbers. •Atomic number and mass number of element X can be denoted as: a<--Mass num Atomic number -->zX<--ELEMENT SYMBOL

The Atomic theory Chapter2

John Dalton was the first scientist to give us a precise definition of the very small, indivisible building blocks of matter that we call atoms.

Radioactivity

•More experimentation led to the discovery of x-rays and then radioactivity. •Radioactivity is the spontaneous emission of particles and/or radiation; any element that spontaneously emits radiation is considered radioactive. Uranium is an example of a radioactive element. Three Types of Rays Produced by Radioactive Decay 1.Alpha () rays: consist of positively charged particles called particles. 2.Beta () rays or particles: are electrons and are deflected by the negatively charged plate. 3.Gamma () rays: high-energy rays, which, like x-rays, have no charge and are not affected by an external electric or magnetic field.

The Electron2

•R. A. Millikan discovered later through experimentation that the charge of an electron is: -1.6022x10^-19 C •Hence, the mass of an electron could be calculated: 𝐦𝐚𝐬𝐬 𝐨𝐟 𝐚𝐧 𝐞𝐥𝐞𝐜𝐭𝐫𝐨𝐧=𝐜𝐡𝐚𝐫𝐠𝐞𝐜𝐡𝐚𝐫𝐠𝐞𝐦𝐚𝐬𝐬 =−𝟏.𝟔𝟎𝟐𝟐×𝟏𝟎^−𝟏𝟗 𝐂/−𝟏.𝟕𝟔×𝟏𝟎^𝟖 𝐂/𝐠 =𝟗.𝟏𝟎×𝟏𝟎^−𝟐𝟖 𝐠

The Proton and the Nucleus2

•Rutherford, working with other scientists, performed -scattering experiments, from which Rutherford proposed a new model of atomic structure. •Rutherford proposed that the atom consisted mostly of empty space. •A structure such as this would allow most of the particles to pass through the gold foil with little or no deflection. •He concluded: 1.The atom's positive charge is concentrated in the nucleus (a dense central core within the atom). 2.The proton (p) has opposite (+) charge of electron (-). 3.The mass of p is 1840 times the mass of e- (1.67 x 10^-24 g).

Guidelines Used When Naming Compounds with Prefixes

•The prefix "mono" may be omitted for the first element. •PCl₃ - Phosphorus trichloride NOT Monophosphorus trichloride •Absence of a prefix means that only 1 atom of that element is present in the molecule. •For oxides, the ending "a" in the prefix is sometimes omitted. •N₂O₄ - Dinitrogen tetroxide NOT Dinitrogen tetraoxide.

Conversion (volume)

𝐕𝐨𝐥𝐮𝐦𝐞 = 𝟏 (𝐦)𝟑 = 𝟏 𝐦𝟑 𝟏 𝐦 = 𝟏𝟎 𝐝𝐞𝐜𝐢𝐦𝐞𝐭𝐞𝐫𝐬 (𝐝𝐦) 𝐕𝐨𝐥𝐮𝐦𝐞 = 𝟏 (𝐦) 𝟑 = (𝟏𝟎 𝐝𝐦)𝟑 = 𝟏𝟎𝟎𝟎 𝐝𝐦𝟑 = 𝟏𝟎𝟎𝟎 𝐋 𝟏 𝐝𝐦 = 𝟏𝟎 𝐜𝐞𝐧𝐭𝐢𝐦𝐞𝐭𝐞𝐫𝐬 (𝐜𝐦) 𝟏 𝐥𝐢𝐭𝐞𝐫 = 𝟏 (𝐝𝐦)𝟑 = (𝟏𝟎 𝐜𝐦)𝟑 = 𝟏𝟎𝟎𝟎 𝐜𝐦𝟑 = 𝟏𝟎𝟎𝟎 𝐦𝐋

Atomic Number, Mass Number, and Isotopes2

Mass Number (A): the total number of neutrons and protons present in the nucleus of an atom of an element. •All atomic nuclei (except for the most common form of hydrogen) contain both protons and neutrons. 𝐌𝐚𝐬𝐬 𝐍𝐮𝐦𝐛𝐞𝐫=𝐍𝐮𝐦𝐛𝐞𝐫 𝐨𝐟 𝐏𝐫𝐨𝐭𝐨𝐧𝐬+𝐍𝐮𝐦𝐛𝐞𝐫 𝐨𝐟 𝐍𝐞𝐮𝐭𝐫𝐨𝐧𝐬=𝐀𝐭𝐨𝐦𝐢𝐜 𝐍𝐮𝐦𝐛𝐞𝐫+𝐍𝐮𝐦𝐛𝐞𝐫 𝐨𝐟 𝐍𝐞𝐮𝐭𝐫𝐨𝐧𝐬 𝐀=𝐙+𝐍𝐮𝐦𝐛𝐞𝐫 𝐨𝐟 𝐍𝐞𝐮𝐭𝐫𝐨𝐧𝐬 𝐍𝐮𝐦𝐛𝐞𝐫 𝐨𝐟 𝐍𝐞𝐮𝐭𝐫𝐨𝐧𝐬=𝐀−𝐙 •Example: If the mass number of a particular boron atom is 12 and the atomic number is 5, then the number of neutrons is 7. •These quantities are always positive integers (whole numbers).

Unit of MM

Mass: Kilogram: kg Length: meter:m Time: second:s Temperature:Kelvin:K

Molecular Models

Molecular Models are used to aid in visualizing molecules. •Two Standard Types of Molecular Models: •Ball-and-Stick Models •Atoms are wooden or plastic balls with holes in them. •Sticks or springs are used to represent chemical bonds. •The angles formed between atoms approximate the bond angles in actual molecules. •With the exception of the H atom, the balls are all the same size and each type of atom is represented by a specific color. •Space-Filling Models: •Represented by truncated balls held together by snap fasteners; therefore, the bonds are not visible. •The balls are proportional in size to atoms. The first step in building a molecular model is to: write the structural formula to show how atoms are bonded to one another in a molecule.

Scientific notation Mult.

Multiplication 𝟒.𝟎×𝟏𝟎−𝟓×𝟕.𝟎×𝟏𝟎𝟑= 1.Combine N1 and N2. 2.Add exponents n1 and n2. 𝟒.𝟎×𝟕.𝟎×𝟏𝟎−𝟓+𝟑= 𝟐𝟖×𝟏𝟎−𝟐=𝟐.𝟖×𝟏𝟎−𝟏

Chemical change

One in which a given substance becomes a new substance or substances with different properties and different composition; alters the composition or identity of the substance (s) involved.

Molecules and Ions - Molecules

Only the 6 noble gases (Group 8A) exist in nature as single atoms. Hence, they are called monatomic (meaning a single atom) gases. • A molecule is an aggregate of at least two atoms in a definite arrangement held together by chemical forces (also called chemical bonds). •Could be atoms of the same element or atoms of two or more elements joined in a fixed ratio according to the law of definite proportions. •A molecule is not necessarily a compound... •Hydrogen gas (H2) is a pure element consisting of molecules made up of two H atoms each. •Water (H2O) is a molecular compound •Molecules are electrically neutral.

Orgo

Organic chemistry is the branch of chemistry that deals with carbon compounds. •Hydrocarbons: the simplest type of organic compounds; contain only carbon and hydrogen atoms. •Used as fuels for domestic and industrial heating •Used for generating electricity and powering internal combustion engines •Starting materials for the chemical industry •Alkanes: one class of hydrocarbons •Straight-chain alkanes: the carbon chains have no branches •All the names end with "-ane" •Starting with C5H12, Greek prefixes are used to indicate the number of carbon atoms present

The Periodic Table

Periodic Table: a chart in which elements having similar chemical and physical properties are grouped together. •Periods: the horizontal rows; elements arranged by atomic number, which is shown above the element symbol). •Groups or Families: the vertical columns; elements arranged according to similarities in chemical properties.

Chemical properties

Properties that do not change the chemical nature of matter. Heat of combustion, reactivity with water, pH, Electromotive force

Physical properties

Properties that do not change the chemical nature of matter.(can be extensive or intensive) Color, odor, freezing point, boiling point and melting point

A&B

Rules for Naming Anions of Oxoacids (Oxoanions) 1.When ALL the H ions are removed from "-ic" acid, the anion's name ends with "-ate": [Example: the anion 𝐂𝐎𝟑𝟐− derived from H2CO3 (carbonic acid) is carbonate] 2.When ALL the H ions are removed from "-ous" acid, the anion's name ends with "-ite": [Example: the anion 𝐂𝐥𝐎𝟐− derived from HClO2 (chlorous acid) is chlorite] 3.The names of anions in which one or more but not all of the hydrogen ions have been removed must indicate the number of H ions present. •Example: Anions derived from phosphoric acid (H3PO4) •𝐇𝟐𝐏𝐎𝟒− − Dihydrogen phosphate •𝐇𝐏𝐎𝟒𝟐− - Hydrogen phosphate •𝐏𝐎𝟒𝟑−Phosphate

Temperature Celsius Vs. Kelvin

Temperature (Kelvin)=Temp. (Celsius)+273.15 Temperature (Celsius)=Temp. (Kelvin)-273.15

Force=mass x acceleration

...

Extensive property

...of a material depends upon how much matter is being considered. *Mass, Length, Volume

Intensive property

...of a material does not depend upon how much matter is being considered. *Density, Temperature, Color