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Honors Chemistry Study Guide
Terms in this set (57)
the law that the properties of the elements are periodic functions of their atomic numbers. 2. Also called Mendeleev's law. (originally) the statement that the chemical and physical properties of the elements recur periodically when the elements are arranged in the order of their atomic weights.
Heat is the form of energy that flows between two samples of matter due to their difference in temperature. Usually denoted by the letter Q or q.
The calorific value is the total energy released as heat when a substance undergoes complete combustion with oxygen under standard conditions. The chemical reaction is typically a hydrocarbon or other organic molecule reacting with oxygen to form carbon dioxide and water and release heat.
Law of Conservation of Energy
The law of conservation of mass states that mass in an isolated system is neither created nor destroyed by chemical reactions or physical transformations. According to the law of conservation of mass, the mass of the products in a chemical reaction must equal the mass of the reactants.
Specific heat is the amount of heat energy required to raise the temperature of a body per unit of mass. In SI units, specific heat (symbol: c) is the amount of heat in joules required to raise 1 gram of a substance 1 Kelvin. Also Known As: specific heat capacity, mass specific heat.
This means that the change in enthalpy under such conditions is the heat absorbed (or released) by the material through a chemical reaction or by external heat transfer. Enthalpies for chemical substances at constant pressure assume standard state: most commonly 1 bar pressure.
Hess's Law of Constant Heat Summation (or just Hess's Law) states that regardless of the multiple stages or steps of a reaction, the total enthalpy change for the reaction is the sum of all changes. This law is a manifestation that enthalpy is a state function.
Heat of Solution
the heat evolved or absorbed when a substance dissolves; specifically : the amount involved when one mole or sometimes one gram dissolves in a large excess of solvent.
a solution in which no more solvent can be dissolved. It is understood that saturation of the solution has been achieved when any additional substance that is added results in a solid precipitate or is let off as a gas.
A solution with solute that dissolves until it is unable to dissolve anymore, leaving the undissolved substances at the bottom. Unsaturated Solution. A solution (with less solute than the saturated solution) that completely dissolves, leaving no remaining substances.
a solution that contains more of the dissolved material than could be dissolved by the solvent under normal circumstances. It can also refer to a vapor of a compound that has a higher (partial) pressure than the vapor pressure of that compound.
In chemistry, a solution is a homogeneous mixture composed of two or more substances. In such a mixture, a solute is a substance dissolved in another substance, known as a solvent.
The component of a solution that is present in the greatest amount. It is the substance in which the solute is dissolved. Examples: The solvent for seawater is water.
theory used to predict the rates of chemical reactions, particularly for gases. The collision theory is based on the assumption that for a reaction to occur it is necessary for the reacting species (atoms or molecules) to come together or collide with one another.
Le Chatelier's Principal
A principle stating that if a constraint (such as a change in pressure, temperature, or concentration of a reactant) is applied to a system in equilibrium, the equilibrium will shift so as to tend to counteract the effect of the constraint.
In a chemical reaction, chemical equilibrium is the state in which both reactants and products are present in concentrations which have no further tendency to change with time. Usually, this state results when the forward reaction proceeds at the same rate as the reverse reaction.
a chemical reaction in which an acid and a base react quantitatively with each other. In a reaction in water, neutralization results in there being no excess of hydrogen or hydroxide ions present in solution.
Arrhenius Acid and Base
An Arrhenius acid is a substance that dissociates in water to form hydrogen ions (H +).
An Arrhenius base is a substance that dissociates in water to form hydroxide (OH -) ions. In other words, a base increases the concentration of OH - ions in an aqueous solution.
Why do atoms form chemical bonds?
They want a full outer shell of electrons, so the lose, gain, or share electrons with other elements, forming compounds, until they have 8 valence electrons and become stable.
A chemical bond formed between two ions with opposite charges. Ionic bonds form when one atom gives up one or more electrons to another atom. These bonds can form between a pair of atoms or between molecules and are the type of bond found in salts.
Soluble in water.
High melting point and boiling point because a large amount of energy is required to break the electrostatic forces holding the lattice together.
They are compounds formed from metals and non-metals.
In a solid state they do not conduct electricity. However, in a liquid state or when dissolved in water, they will conduct electricity well because the ions are free to move.
A covalent bond, also called a molecular bond, is a chemical bond that involves the sharing of electron pairs between atoms. These electron pairs are known as shared pairs or bonding pairs, and the stable balance of attractive and repulsive forces between atoms, when they share electrons, is known as covalent bonding.
Low melting points and boiling points.
Low enthalpies of fusion and vaporization These properties are usually one or two orders of magnitude smaller than they are for ionic compounds.
Soft or brittle solid forms.
Poor electrical and thermal conductivity.
Metallic bonding is the principal force holding together the atoms of a metal. A metallic bond results from the sharing of a variable number of electrons by a variable number of atoms.
High thermal and electrical conductivity.
Luster and high reflectivity.
Malleability and ductility. They can be beaten or shaped without fracture.
Variability of mechanical strengths (ranging from soft alkali metals to Tungsten, which is hard).
Contrast single, double, and triple covalent bonds
Single bond - sharing of 1 electron between elements. Example: H2, Cl2, H2O
Double bond - sharing of 2 electrons between elements. Example: O2, CO2
Triple bond - sharing of 3 electrons between elements. Example: N2
London Dispersion Forces
These are the weakest of the intermolecular forces and exist between all types of molecules, whether ionic or covalent—polar or nonpolar. The more electrons a molecule has, the stronger the London dispersion forces are.
These forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule. The prerequisite for this type of attraction to exist is partially charged ions—for example, the case of polar covalent bonds such as hydrogen chloride. Dipole-dipole interactions are the strongest intermolecular force of attraction.
This is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen, nitrogen, or fluorine atom. The partially positive end of hydrogen is attracted to the partially negative end of the oxygen, nitrogen, or fluorine of another molecule. Hydrogen bonding is a relatively strong force of attraction between molecules, and considerable energy is required to break hydrogen bonds. This explains the exceptionally high boiling points and melting points of compounds like water, and hydrogen fluoride.
How to draw a Lewis Structure
Step 1: Find the Total Number of Valence Electrons.
Step 2: Find the Number of Electrons Needed to Make the Atoms "Happy".
Step 4: Choose a Central Atom.
Step 5: Draw a Skeletal Structure.
Step 6: Place Electrons Around Outside Atoms.
STEP 7: Place remaining electrons around the central atom.
Trigonal Planar 120
Bent Less Than 120
A solid is a state of matter characterized by particles arranged such that their shape and volume are relatively stable. The constituents of a solid tend to be packed together much closer than the particles in a gas or liquid.
An amorphous (non-crystalline) form of matter between a gas and a solid that has a definite volume, but no definite shape.
One of four main states of matter, composed of molecules in constant random motion. Unlike a solid, a gas has no fixed shape and will take on the shape of the space available. Unlike a liquid, the intermolecular forces are very small; it has no fixed volume and will expand to fill the space available.
What increases the pressure of a gas?
An increase in the number of gas molecules in the same volume container increases pressure. A decrease in container volume increases gas pressure. An increase in temperature of a gas in a rigid container increases the pressure.
When the system is heated, energy is transferred into it. In response to the energy it receives, the system changes, for example by increasing its temperature. A plot of the temperature versus time is called the heating curve.
A diagram representing the limits of stability of the various phases in a chemical system at equilibrium, with respect to variables such as composition and temperature.
A graphic representation of the variation with changing temperature of the solubility of a given substance in a given solvent.
Endothermic and Exothermic Reaction
An exothermic reaction occurs when the temperature of a system increases due to the evolution of heat. This heat is released into the surroundings, resulting in an overall negative quantity for the heat of reaction.
An endothermic reaction occurs when the temperature of an isolated system decreases while the surroundings of a non-isolated system gains heat. Endothermic reactions result in an overall positive heat of reaction.
Exothermic and endothermic reactions cause energy level differences and therefore differences in enthalpy (ΔH), the sum of all potential and kinetic energies. ΔH is determined by the system, not the surrounding environment in a reaction. A system that releases heat to the surroundings, an exothermic reaction, has a negative ΔH by convention, because the enthalpy of the products is lower than the enthalpy of the reactants of the system.
Conductors vs Insulators
As heat is added uniformly to like quantities of different substances, their temperatures can rise at different rates. For example, metals, good conductors of heat, show fast temperature rises when heated. It is relatively easy to heat a metal until it glows red. On the other hand, water can absorb a lot of heat with a relatively small rise in temperature. Insulating materials (insulators) are very poor conductors of heat, and are used to isolate materials that need to be kept at different temperatures — like the inside of your house from the outside.
Calorimetry is the measurement of the quantity of heat exchanged. For example, if the energy from an exothermic chemical reaction is absorbed in a container of water, the change in temperature of the water provides a measure of the amount of heat added.
A calorimeter is an apparatus for measuring the amount of heat involved in a chemical reaction or other process.
Why do electrolytes affect colligative properties differently than non-electrolytes?
Electrolytes produce more moles of solute particles per mole of solute.
Net Ionic Equations
Net ionic equations are equations that show only the soluble, strong electrolytes reacting (these are represented as ions) and omit the spectator ions, which go through the reaction unchanged.
Experimental Determination of Rate Law. The values of k, m, and n in the rate law equation must be determined experimentally for a given reaction at a given temperature. The rate is usually measured as a function of the initial concentrations of the reactants, A and B.
Factors affecting reaction rates
surface area of a solid reactant
concentration or pressure of a reactant
nature of the reactants
presence/absence of a catalyst.
Assumptions about the Reaction Based on the Value of K. When we know the numerical value of the equilibrium constant, we can make certain judgments about the extent of the chemical reaction. If K is larger than 1, the mixture contains mostly products. If K is less than 1, the mixture contains mostly reactants.
HCl - hydrochloric acid.
HNO3 - nitric acid.
H2SO4 - sulfuric acid (HSO4- is a weak acid)
HBr - hydrobromic acid.
HI - hydroiodic acid.
HClO4 - perchloric acid.
HClO3 - chloric acid.
LiOH - lithium hydroxide.
NaOH - sodium hydroxide.
KOH - potassium hydroxide.
RbOH - rubidium hydroxide.
CsOH - cesium hydroxide.
*Ca(OH)2 - calcium hydroxide.
*Sr(OH)2 - strontium hydroxide.
Acids vs Bases
pH < 7.
Sour taste (though you should never use this characteristic to identify an acid in the lab)
Reacts with a metal to form hydrogen gas.
Increases the H+ concentration in water.
Donates H+ ions.
Turns blue litmus indicator red.
(H+ is short for H3O+)
pH > 7
Increases the OH- concentration in water
Accepts OH- ions
Turns red litmus indicator blue
Conjugate Acids and Bases
A conjugate acid, within the Brønsted-Lowry acid-base theory, is a species formed by the reception of a proton (H+) by a base—in other words, it is a base with a hydrogen ion added to it. On the other hand, a conjugate base is merely what is left after an acid has donated a proton in a chemical reaction.
Almost all stoichiometric problems can be solved in just four simple steps:
Balance the equation.
Convert units of a given substance to moles.
Using the mole ratio, calculate the moles of substance yielded by the reaction.
Convert moles of wanted substance to desired units.
Dalton's Law of Partial Pressure
Dalton's law (also called Dalton's law of partial pressures) states that in a mixture of non-reacting gases, the total pressure exerted is equal to the sum of the partial pressures of the individual gases.
Combined Gas Law
The combined gas law combines the three gas laws: Boyle's Law, Charles' Law, and Gay-Lussac's Law. It states the ratio of the product of pressure and volume and the absolute temperature of a gas is equal to a constant.
The specific heat capacity of a substance is the amount of heat required to raise one gram of the substance by one degree Celsius. Water, for example, has a specific heat capacity of 4.18 . This means to heat one gram of water by one degree Celsius, it would require 4.18 joules of energy.
You can solve for the concentration or volume of the concentrated or dilute solution using the equation: M1V1 = M2V2, where M1 is the concentration in molarity (moles/Liters) of the concentrated solution, V2 is the volume of the concentrated solution, M2 is the concentration in molarity of the dilute solution
Molarity is a concentration unit, defined to be the number of moles of solute divided by the number of liters of solution.
a number that expresses the relationship between the amounts of products and reactants present at equilibrium in a reversible chemical reaction at a given temperature.
a measure of hydrogen ion concentration; a measure of the acidity or alkalinity of a solution. Aqueous solutions at 25°C with a pH less than seven are acidic, while those with a pH greater than seven are basic or alkaline.
the slow addition of one solution of a known concentration (called a titrant) to a known volume of another solution of unknown concentration until the reaction reaches neutralization, which is often indicated by a color change.
Ka values with ICE Charts
Use the calculator
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