Chemistry exam 1

Chapter 1

Science and Measurement 

Aug 21, 2024 | study

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Notes 1.1:

Chemistry is the study of matter and energy

Matter is anything that has mass and occupies space 

 Elements are the simplest form of matter that has distinct physical and chemical properties (discussed in Section 1.2) and cannot be broken down chemically into simpler, stable substances. They are the building blocks for everything in the universe.

An atom is the smallest amount of an element that still has the characteristic of that element. Atoms of different elements can form attractions called Chemical Bonds. When atoms create chemical bonds they form compounds. 

Compounds are a chemical combination of elements that has it’s own set of properties and a definite composition. For example, a definite composition could be water, 2 hydrogens, and 1 oxygen.

Elements and compounds are pure substances. 

A Mixture is two or more substances that are physically combined. They are not chemically bonded. They can be separated by physical means, for example, filtering. Mixtures do not have definite compositions. 

There are two types of Mixtures. First is a heterogeneous mixture, which has substances that are not uniformly mixed. Next is a homogeneous mixture which has a uniform composition, they are also called a solution.

Metal alloy can be called a solid solution, salt water can be called an aqueous solution.

Notes 1.2:

Every substance that we have has its own set of properties, which are characteristics by which something can be identified. A physical property is something to describes a substance without changing its chemical composition, for example, color. Chemical properties are characteristic chemical reactions a substance undergoes, for example, rusting.

Iron has the physical property of being a gray solid at room temperature. It also has a chemical property which is iron reacts with oxygen in the air to form rust.

Extensive properties depend on the amount of substance present like mass or volume.

Intensive properties are the same regardless of sample size.

Properties of compounds are: constant and typically different from the properties of elements that compose them. For example, sodium and table salt. 

Properties of mixtures are: Similar to properties of the components that make up the mixture. They also change depending on the amount of each component. For example, sugar water.

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Aug 23, 2024 | study

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Notes 1.2:

Physical change: A physical change is when the chemical composition of a substance is not altered. For example, when mixing iron and sulfur they remain iron and sulfur so in other words, there is no chemical reaction. Another example is when carbon dioxide (dry ice) goes into a gas phase, but its chemical composition is the same (CO2).

Chemical change: A chemical change is when a starting material, a reactant(s), forms a product, or resulting material. The reactants rearrange their chemical compositions to form products. For example, when iron and sulfur are heated together a new substance forms with different properties than iron, or sulfur.

Chemical Change:

Changes the chemical structure of a substance by:

Breaking and/or forming chemical bonds.

When bonds are broken:

Energy is absorbed from the surroundings.

When bonds are formed:

Energy is released to the surroundings.

Notes 1.3:

Mass of an object: measures how much matter is in an object

Weight of an object: measures the heaviness (the force of gravity on the object) of an object,

Mass and weight are: directly proportional

*Many scientists use mass and weight interchangeably however it only applies when working on Earth’s surface.

Energy: the capacity to do work.

Law of conservation of energy: states that energy cannot be created or destroyed but can be converted from one form to another.

Examples of forms of energy: Heat, kinetic energy, potential energy, chemical energy, 

1.4 The Scientific Method:

The scientific method is the process of conducting experimental science. 

Observation: background information/data

Law: Observations that are always true. Because they are always true they are statements without explanation.

Hypothesis: Initial explanation for the background info/data. Tested and revised through experimentation.

Experiment: object the hypothesis

Scientific theory: This is something that explains/predicts different observations linked by the underlying phenomenon. Generally excepted as a valid explanation. 

Law of conservation of mass: says that in any chemical reaction or physical change, the total mass present after the change is equal to the total mass present before the change.

• In the 20th century, this was found to not apply to nuclear reactions.

Notes 1.5:

The International system of units (SI units) is a system of units based on the metric system, which includes units like meters, kilograms, and seconds.

Base units are defined by a particular physical measurement: 

SI Prefixes: added to a unit to describe a very large or very small measurement.

For example, a kilogram has an abbreviation of k, and it tells us that the grams are multiplied by 10^3

SI derived units: Area and Volume. They are the product of one or more base units.

5N = 5kg*m/s^2

SI English conversions:

Notes 1.6: Significant figues

Qualitative Data: refers to the identity or form of a substance

Quantitative Data: determine the amount of a substance.

Scientific measurements are usually repeated several times because the average value is probably closer to the true value than any individual measurement. 

Precision: closeness of all of a set of measurements to one another

Precision in single measurements is determined by the measuring device used.

When making measurements: estimate to one digit beyond the smallest scale division on the tool. For example, a ruler measures to the tenth place, so estimate to the hundredth place.

Accuracy: The closeness of the average of a set of measurements to the true value. Accuracy means precision, but precision does not mean accuracy.

Scientists report the precision of their measurements every time they write down a result.

Rules for significant figures:

Significant figures (sig figs) are digits that reflect precision. They can be defined as absolutely certain digits plus one estimated digit. 

All nonzero numbers are significant

Any zeros between sigfig digits are significant

Any zeros to the right of all nonzero digits in a number with decimal place digits are significant

All digits in the coefficient of a number in scientific notation are significant.

Any zeros to the right of all nonzero digits in an integer are not significant

Any zeros to the left of all nonzero digits are not significant

Significant figures in addition and subtraction the answer is rounded to the fewest number of decimal places present in the values

Significant figures in multiplication and division the answer is rounded to the fewest number of significant figures present in the values

Exact numbers: Numbers that are definitions(not measurements). For example 100cm/1m. 

  Counted items. For example students in our classroom

  Integers in formulas. d = 2r

Exact numbers do not determine the significant figures in a calculated result

Sigfigs with mixed operations:

Follow the order of operations (PEMDAS). The part done first must have its Sigfigs noted before the next operation is performed. To avoid round off errors retain at least one extra digit until the final answer.

Notes 1.7 dimensional analysis:

Every measurement has a number and a unit. 

Dimensional analysis involves the use of units to do calculations. Often include a conversion factor.

A conversion factor is a ratio equal to one that can be multiplied by a quantity to change the form of the quantity without changing its value.

Notes 1.8 Density:

Density: The measure of how much mass something has relative to how much space it takes up.

Density = mass/volume or d=m/v

Units: Both mass and volume units. For example g/mL or g/cm^3

It can be used as a conversion factor between mass and volume

Density is an intensive property or one that is independent of size.

Density of water: 1.0 g/mL

1.9 Temperature scales:

Fahrenheit: boiling point = 212 degrees. Freezing point = 32 degrees

Celcius: boiling point = 100 degrees. Freezing point = 0 degrees

Kelvin: SI Unit for chemistry. Boiling point = 373.15 k. Freezing point = 273.15 k. The coldest possible temperature in Kelvin is 0 k. This is called absolute zero. It is equal to -459.67 degrees F or -273.15 degrees Celsius.

Converting between temperatures:

Sep 4, 2024 | Chem 105 Fulmer Hall 0226

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Notes: Atoms and the Periodic Table 2.4

Subatomic particles: include protons, neutrons, and electrons

Nucleus: Almost all of the mass of an atom. A nucleus has a radius of about one ten-thousandth the size of the atom's radius. Electrons make up the most of the volume of an atom.

Atoms are electronically neutral because their number of protons is equal to their number of electrons. (only true for a neutral atom)

Neutrons: are electrically neutral, which means they do not affect the charge of the atom.

Atomic number (z) is equal to (p) or, the number of protons in the atom's nucleus. The atomic number determines the atom's identity because it’s distinct for each element.

Number of neutrons: can differ in atoms of the same element.

Isotopes: same number of protons and a different number of neutrons. They are the same element because they have the same atomic number. However, they have different masses because they have a different number of neutrons.

Mass number (A): is our number of protons plus our number of neutrons. Or A = p + n, or A = z + n

Symbols for isotopes: 

Hydrogen - 1 or ₁¹H one on top = A = p + n and one on bottom = Z = p.

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Ions: Atoms that are not neutral. They are charged particles formed when atoms gain or lose electrons. 

Atoms can react to lose an electron and form a positively charged ion called a cation.

Atoms can react to gain an electron and form a negatively charged ion called an anion.

Notes 2.5:

Atomic mass scale: a relative scale with one isotope assigned a value and all others measured relative to that isotope. Our standard in which all other isotopes are measured against is carbon - 12 which has exactly 12 atomic mass units (u).

An atomic mass unit is so small that it takes 6.027 x 10²³ u to make 1.00g.

Determination of atomic mass: we use the weighted average of the actual masses of the element’s naturally occurring isotopes on Earth.

Atomic mass = ∑ fraction of isotope x mass of isotope

Atomic mass is from a mixture of naturally occurring isotopes. The mass number refers to a single isotope.

Artificial elements: are elements with their atomic number (Z) greater than 92. Artificial elements do not have atomic masses because there are no naturally occurring isotopes. On the periodic table, the number is the mass number of the most stable isotope, its usually written in parentheses. Artificial isotopes do not contribute to the atomic mass. 

Notes 2.6:

History of the periodic table: Initially, elements were ordered by atomic mass. Took many years to refine the periodic table. The modern periodic table has elements arranged horizontally by increasing atomic number. They are also grouped vertically by similar properties. 

Elements within a horizontal row are in the same period. 

The 6th and 7th periods include the inner transition metals. (lanthanoids and actinoids)

Elements within a vertical column are in the same group or family. Elements in the same group or family have similar chemical properties. We number them from left to right. The old system is 1-8 with A or B and the new system is 1-18.

Group one is called the alkali metals. Group two is alkaline earth metals. 11 is coinage metals. 17 is halogens. 18 is noble gases. 

The majority of elements are metals. Metals are shiny, malleable, and ductile. Malleable means it can be pounded into thin sheets and ductile means it can be drawn into a wire. Metals are also conductors of heat and electricity.

Nonmetal elements are a dull color and typically brittle as solids. They are not conductors of head or electricity. 

Metalloid elements share properties with both nonmetals and metals. 

Sep 9, 2024 | Chem 105 Fulmer Hall 0226

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Notes 3.1 Chemical Formulas

Chemical Nomenclature: a systematic approach for naming compounds

Chemical Formulas: combines element symbols (to represent the element) and subscripts (to represent the amount of the element) to represent compounds

Chemical formulas provide information about the relative number of atoms of each element in a compound. It can also provide information about how atoms are arranged in the compound.

Formula unit: the group of atoms represented by the chemical formula

The formula unit represents: The ratio of atoms in a compound

Molecules: individual particles of the compound. Molecular compounds contain nonmetal-nonmetal bonds, or nonmetal-metalloids. We refer to these as covalent bonds.

Ionic compounds exist as extended 3D lattice structures (repeating units or patterns) of atoms. There are no distinct molecules in an ionic compound. Instead, the formula unit represents repeating patterns of bonded atoms. Atoms are held together by ionic bonds. (Chp 9). Ionic compounds contain metal-nonmetal bonds.

Some elements are never found in nature as individual atoms.

Diatomic molecules: are molecules consisting of two atoms. H2, N2, O2, F2, Cl2, Br2, I2

Allotropes: Different molecular forms of the same element.

Notes 3.2 Name binary covalent compounds

Binary compounds: Made up of two elements. 

Binary covalent compound: made up of two nonmetals.

The same two elements can bond together in different ratios. 

Name using prefixes to indicate the number of each type present. 

Binary compounds of Hydrogen: Many are acids and are named using the rules for naming acides

Those that aren’t acids have special names (memorize):

NH3: ammonia

PH3: phosphine

AsH3: Arsine

Notes 3.3:

Monatomic ions: single atoms that gain/lose electrons. Don’t have multiple types of ions.

Binary ionic compounds: form when metal atoms combine with nonmetal atoms. The metal atom transfers electrons to the nonmetal atoms. Metal atoms form cations. Nonmetal atoms form an anion. 

Compounds are electrically neutral overall. Total positive charge = total negative charge

Polyatomic Ions: groups of two or more bonded atoms that have lost or gained electrons. Most are anions

3.4: Naming Ionic compounds 

The only common polyatomic cation is: NH4 +

Metal cation with constant charge: name of element + ion

Metal cation with variable charge: name + Roman numeral + ion

Naming monatomic anions: all monatomic anions have a constant charge. They form from groups 15-17. To name them we use the name of the element + change the ending to “ide” + “ion”.

Naming polyatomic anions by memorizing the charges of the polyatomic anions. 

Oxyanions: contain one nonmetal + variable number of oxygens 

The formula of an ionic compound indicates the lowest whole number ratio of cations to anions. No need to include ratio information

Hydrates: Ionic compounds that contain water molecules within their solid structure 

A hydrate formula shows waters of hydration

Naming hydrates: add a prefix and add hydrate to the end of the name

Heating a hydrate: would remove the water molecules and create its anhydrous form 

(without water)

Notes 3.5 Naming acids

Acids: special class of H-containing covalent compounds 

When dissolved in water: releases 1 or more H+ ions and forms an anion with a 1- charge for each H+ ion that's released.

Ionizable hydrogen atoms are the H+ ions released in water. They are written at the beginning of a formula 

Two types of acids: binary acids and oxyacids

Monoprotic acids: 1 ionizable hydrogen

Polyprotic acids: more then 1 ionizable hydrogen 

Only hydrogen atoms in the front of an acid can be ionized.

Binary acids: contain hydrogen and an atom of an element from group 16 or 17

The number of H atoms = the # of negative charges on the anion

Oxyacids: H atoms bonded to oxyanions

3.7 The mole

1 dozen = 12 units of anything 

Chemical counting unit: Mole (mol). It’s used to count atoms, molecules, and formula units

1 mole of anything: 6.022 * 10^23 units of that thing

Avagadros number (N_A): 6.022 * 10^23

Moles of atoms in a compound 

Use the formula to obtain the mole ratios within the compound

Determine atoms in a compound:

If you know the number of moles of an element you can use Avagodros number to 

determine the number of atoms.

3.8 molar mass:

Review: the atomic mass (u or amu) of each element is the average mass of one atom

Atomic mass units in a gram: 1g/6.022x10^23 u

Mass of 1 mole of atoms: the atomic mass expressed in grams 

For example (6.022x10^23 g C) (12.011 u C/1 C atom) (1g/6.022x10^23 u C) = 

12.011 g C

The atomic mass of C = 12.011 u. Molarmass of C is equal to 12.011 grams per mole.

The molar mass of compounds is the sum of the atomic mass of all the elements/atoms in a formula unit of the compound.

Atomic mass is the average mass of an atom, expressed in u

Molecular mass is the average mass of a molecule expressed in u

Formula mass is the average mass of a formula unit of any type of substance (atom, molecule, or ionic compound), expressed in u