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enthalpy change of reactions
calorimetry
Q=mcΔT
when data of calorimetry experiment is absent
if two or more other reactions and their enthalpy changes are given
use Hess's law
if bond enthalpies of reactants are products are given
ΔH= sum of bond enthalpies of reactants- sum of bond enthalpies of products
if molar enthalpies of formation of reactants and products are given
ΔH= ∑molar enthalpy of formation of products - ∑molar enthalpy of formation of reactants
solid
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Periodic Table
elements
All you can see on periodic table. In chemistry, any material (such as carbon, hydrogen, iron, or oxygen) that cannot be broken down into more fundamental substances. Each chemical element has a specific type of atom, and chemical compounds are created when atoms of different elements are bound together into molecules. There are 119 chemical elements whose discovery has been claimed; 92 occur in nature, and the rest have been produced in laboratories.
Chemical changes in compounds happen when chemical bonds are created or destroyed. Forces act on the bonds between atoms, changing the molecular structure of a substance. You can pour liquid acid on a solid and watch the solid dissolve. That process is a chemical change because molecular bonds are being created and destroyed. Geologists pour acids on rocks to test for certain compounds.
atoms
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Structure of a beryllium atom: four protons, four neutrons and four electrons.
Structure of a beryllium atom: four protons, four neutrons and four electrons.
Credit: general-fmv Shutterstock
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Atoms are the basic units of matter and the defining structure of elements. The term "atom" comes from the Greek word for indivisible, because it was once thought that atoms were the smallest things in the universe and could not be divided. We now know that atoms are made up of three particles: protons, neutrons and electrons — which are composed of even smaller particles such as quarks.
molecules
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A molecule is the smallest particle in a chemical element or compound that has the chemical properties of that element or compound. Molecules are made up of atoms that are held together by chemical bonds. These bonds form as a result of the sharing or exchange of electrons among atoms. The atoms of certain elements readily bond with other atoms to form molecules. Examples of such elements are oxygen and chlorine. The atoms of some elements do not easily bond with other atoms. Examples are neon and argon.
compounds
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A compound is a molecule made of atoms from different elements. All compounds are molecules, but not all molecules are compounds. Hydrogen gas (H2) is a molecule, but not a compound because it is made of only one element. Water (H2O) can be called a molecule or a compound because it is made of hydrogen (H) and oxygen (O) atoms.
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Properties of Water
mixtures
homogenous
heterogenous
solutions
solvent
solute
Bohr's-Rutherford Diagram
An ion is a positively or negatively charged particle and can also be an atom or molecule that has lost or gained one or more electrons. An ion itself primarily consists of protons, neutrons and electrons. The composition of these particles can describe the type of atom that has been created. For example, an atom with an equal number of protons and neutrons is considered neutral, whereas an atom that has more protons than electrons is considered a positively charged atom. On the other hand, atoms filled with more electrons are considered negatively charged.
Lewis dot Diagram
Creative Chemistry Hangman Game
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Atomic Number
Atomic Mass
Ionic and Covalent Boding
SubThere are two main types of chemical bonds that hold atoms together: covalent and ionic/electrovalent bonds. Atoms that share electrons in a chemical bond have covalent bonds. An oxygen molecule (O2) is a good example of a molecule with a covalent bond. Ionic bonds occur when electrons are donated from one atom to another. Table salt (NaCl) is a common example of a compound with an ionic bond.
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What are isotopes
WHAT IS AN ISOTOPE
An isotope is similar to an ion with one fundamental difference. While both ions and isotopes have a specific number of protons and electrons, an isotopic atom is specific to a particular element, like say, hydrogen. The number of protons and electrons are also often atypical, or not normal for a basic atom. This means ions are easily defined by the exact number of protons and neutrons in the nucleus. This tells you what kind of element it is.
Examples of Isotopes
Isotopes can also be a variation of an element. An example of this would be the difference between hydrogen and deuterium, which is a form of hydrogen. Pure hydrogen is devoid of a neutron while deuterium has two. Both of these elements have isotopic atoms.
atomic radius
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Kinematics
Different Dimensions
1D
Simple addition and subtraction
2D
Must involve vectors with directions
May require using trigonometry and Sine Law
Acceleration and velocities
Force of gravity (9.8m/s2)
Initial and Final Velocity
Equations
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Naming Compounds
Chemical properties
explanation
Physical properties
Sink or Float with the Cookie Monster
5 senses
Types of Matter
liquid
gases
Solid
Electronegativity
What is electronegativity
Definition
Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. Electronegativity differs from electron affinity because electron affinity is the actual energy released when an atom gains an electron. Electronegativity is not measured in energy units, but is rather a relative scale. All elements are compared to one another, with the most electronegative element, fluorine, being assigned an electronegativity value of 3.98. Fluorine attracts electrons better than any other element. The table below shows the electronegativity values for the elements.
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ionization energy
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Ionization energy is the energy required to remove an electron from a gaseous atom or ion. The first or initial ionization energy or Ei of an atom or molecule is the energy required to remove one mole of electrons from one mole of isolated gaseous atoms or ions.
You may think of ionization energy as a measure of the difficulty of removing electron or the strength by which an electron is bound. The higher the ionization energy, the more difficult it is to remove an electron. Therefore, ionization energy is in indicator of reactivity. Ionization energy is important because it can be used to help predict the strength of chemical bonds.
Also Known As: ionization potential, IE, IP, ΔH°
Units: Ionization energy is reported in units of kilojoule per mole (kJ/mol) or electron volts (eV).
Ionization Energy Trend in the Periodic Table
Ionization, together with atomic and ionic radius, electronegativity, electron affinity, and metallicity, follows a trend on the periodic table of elements.
Ionization energy generally increases moving from left to right across an element period (row). This is because the atomic radius generally decreases moving across a period, so there is a greater effective attraction between the negatively charged electrons and positively-charged nucleus. Ionization is at its minimum value for the alkali metal on the left side of the table and a maximum for the noble gas on the far right side of a period. The noble gas has a filled valence shell, so it resists electron removal.
Ionization decreases moving top to bottom down an element group (column). This is because the principal quantum number of the outermost electron increases moving down a group. There are more protons in atoms moving down a group (greater positive charge), yet the effect is to pull in the electron shells, making them smaller and screening outer electrons from the attractive force of the nucleus. More electron shells are added moving down a group, so the outermost electron becomes increasingly distance from the nucleus.
First, Second, and Subsequent Ionization Energies
The energy required to remove the outermost valence electron from a neutral atom is the first ionization energy. The second ionization energy is that required to remove the next electron, and so on. The second ionization energy is always higher than the first ionization energy. Take, for example, an alkali metal atom. Removing the first electron is relatively easy because its loss gives the atom a stable electron shell. Removing the second electron involves a new electron shell that is closer and more tightly bound to the atomic nucleus.
The first ionization energy of hydrogen may be represented by the following equation:
ΔH° = -1312.0 kJ/mol
Exceptions to the Ionization Energy Trend
H(g) → H+(g) + e-
If you look at a chart of first ionization energies, two exceptions to the trend are readily apparent. The first ionization energy of boron is less than that of beryllium and the first ionization energy of oxygen is less than that of nitrogen.
The reason for the discrepancy is due to the electron configuration of these elements and Hund's rule. For beryllium, the first ionization potential electron comes from the 2s orbital, although ionization of boron involves a 2p electron. For both nitrogen and oxygen, the electron comes from the 2p orbital, but the spin is the same for all 2p nitrogen electrons, while there is a set of paired electrons in one of the 2p oxygen orbitals.
Key Points
Ionization energy is the minimum energy required to remove an electron from an atom or ion in the gas phase.
The most common units of ionization energy are kilojoules per mole (kJ/M) or electron volts (eV).
Ionization energy exhibits periodicity on the periodic table.
The general trend is for ionization energy to increase moving from left to right across an element period. Moving left to right across a period, atomic radius decreases, so electrons are more attracted to the (closer) nucleus.
The general trend is for ionization energy to decrease moving from top to bottom down a periodic table group. Moving down a group, a valence shell is added. The outermost electrons are further from the positive-charged nucleus, so they are easier to remove.