Cheat Sheet
Atomic Theory Cheat Sheet
1. Dalton's Atomic Theory (1803):
Proposed by John Dalton.
Elements are composed of indivisible particles called atoms.
All atoms of the same element are identical in mass and properties.
Compounds are formed by the combination of atoms in simple, whole-number ratios.
Chemical reactions involve the rearrangement of atoms, and atoms are neither created nor destroyed.
2. Modern Atomic Theory:
Atoms are the basic building blocks of matter.
Atoms are composed of subatomic particles:
Protons (positively charged) in the nucleus.
Neutrons (neutral) in the nucleus.
Electrons (negatively charged) orbiting the nucleus in electron clouds or energy levels.
Elements are defined by the number of protons, known as the atomic number.
3. Subatomic Particles:
Protons:
Mass: ~1 atomic mass unit (amu).
Charge: +1.
Neutrons:
Mass: ~1 amu.
Charge: 0.
Electrons:
Mass: ~1/1836 amu (negligible).
Charge: -1.
4. Atomic Structure:
Atomic Number (Z): The number of protons in an atom.
Mass Number (A): The sum of protons and neutrons in the nucleus.
Isotopes: Atoms of the same element with different numbers of neutrons.
Atomic Mass: Weighted average of the masses of all naturally occurring isotopes of an element.
5. The Periodic Table:
Organizes elements based on atomic number and chemical properties.
Groups (columns) have similar chemical properties.
Periods (rows) represent the number of electron shells.
6. Atomic Models:
Thomson's Model (Plum Pudding Model): Electrons embedded in a positively charged sphere.
Rutherford's Model: Discovered the atomic nucleus, where most of the mass is concentrated.
Bohr's Model: Electrons orbit the nucleus in discrete energy levels or shells.
Modern Quantum Mechanical Model: Describes the behavior of electrons as waves and particles in electron clouds or orbitals.
7. Quantum Numbers:
Principal Quantum Number (n): Represents the energy level or shell (n=1, 2, 3, ...).
Angular Momentum Quantum Number (l): Determines the shape of the orbital (s, p, d, f).
Magnetic Quantum Number (m_l): Specifies the orientation of an orbital in space.
Spin Quantum Number (m_s): Describes the spin of an electron (up or down).
8. Electron Configuration:
The arrangement of electrons in energy levels and orbitals.
Follows the Aufbau principle, Pauli exclusion principle, and Hund's rule.
9. Periodic Trends:
Atomic radius: Increases down a group, decreases across a period.
Ionization energy: Decreases down a group, increases across a period.
Electronegativity: Decreases down a group, increases across a period.
10. Chemical Bonding:
Atoms form bonds to achieve a stable electron configuration.
Types of bonds include covalent, ionic, and metallic.
11. Molecules and Compounds:
Molecules are formed when atoms share electrons (covalent bonding).
Compounds are substances composed of different elements in fixed ratios.
12. Nuclear Reactions:
Involves changes in the nucleus.
Examples include nuclear decay, fusion, and fission.
13. Modern Advances:
The discovery of subatomic particles beyond protons, neutrons, and electrons, such as quarks and neutrinos.
The development of quantum mechanics and its application to atomic theory.
Structure and Bonding Cheat Sheet
1. Types of Chemical Bonds:
Covalent Bonds:
Formed by the sharing of electrons between atoms.
Can be polar (unequal sharing) or nonpolar (equal sharing).
Examples: H2 (hydrogen gas), O2 (oxygen gas).
Ionic Bonds:
Formed when electrons are transferred from one atom to another, resulting in ions.
Ions are held together by electrostatic attractions.
Example: NaCl (sodium chloride).
2. Lewis Dot Structures:
Diagrams representing the valence electrons of atoms.
Useful for predicting and understanding bond formation.
Follows the octet rule: Atoms tend to gain, lose, or share electrons to achieve a full outer shell of 8 electrons (2 for hydrogen).
3. VSEPR Theory (Valence Shell Electron Pair Repulsion):
Predicts molecular geometry based on the arrangement of valence electron pairs around the central atom.
Minimizes repulsion between electron pairs.
Helps determine molecular shape (e.g., linear, bent, tetrahedral).
4. Intermolecular Forces:
Forces of attraction between molecules.
Types:
London Dispersion Forces: Present in all molecules; temporary fluctuations in electron distribution.
Dipole-Dipole Interactions: Occur in polar molecules due to uneven electron distribution.
Hydrogen Bonds: A special type of dipole-dipole interaction involving hydrogen bonded to highly electronegative elements (N, O, F).
5. Valence Bond Theory:
Describes chemical bonding in terms of overlapping atomic orbitals.
Covalent bonds are formed when orbitals of two atoms overlap.
Examples include sigma (σ) and pi (π) bonds.
6. Molecular Orbital Theory:
Describes the behavior of electrons in terms of molecular orbitals (MOs).
MOs are formed by the linear combination of atomic orbitals (LCAO).
Predicts bond order, bond strength, and magnetic properties.
Examples include bonding, antibonding, and nonbonding orbitals.
7. Ionic Compounds:
Composed of cations (positively charged ions) and anions (negatively charged ions).
Held together by electrostatic attractions.
Form crystal lattices.
Conduct electricity when molten or dissolved in water.
8. Covalent Compounds:
Formed by the sharing of electrons.
Can be polar or nonpolar based on electronegativity differences.
Often exist as discrete molecules.
Generally have lower melting and boiling points compared to ionic compounds.
9. Hybridization:
The process by which atomic orbitals mix to form new hybrid orbitals.
Common in organic compounds and explains molecular geometry.
10. Bond Polarity vs. Molecular Polarity:
Bond polarity is determined by the electronegativity difference between bonded atoms.
Molecular polarity depends on the shape and polarity of the bonds within a molecule.
Transition Metals Cheat Sheet
1. Transition Metals:
Located in groups 3 to 12 on the periodic table.
Also known as the "d-block" elements.
Exhibit a wide range of properties due to their partially filled d-orbitals.
Often form colorful compounds and serve as catalysts.
2. Electron Configurations of Transition Metals:
Transition metals have electron configurations that include d-orbitals in addition to s and p orbitals.
The filling order follows the Aufbau principle, Hund's rule, and the Pauli exclusion principle.
For example, the electron configuration of chromium (Cr) is [Ar] 3d^5 4s^1, and not [Ar] 3d^4 4s^2 as might be expected.
3. Properties of Transition Metals:
Variable Oxidation States:
Transition metals can exist in multiple oxidation states due to the availability of d-electrons.
Some elements have particularly stable oxidation states, such as Fe(II) and Fe(III) for iron.
Complex Formation:
Transition metals readily form coordination complexes with ligands (molecules or ions).
Ligands coordinate to the central metal ion through dative covalent bonds.
Complexes often exhibit color due to electronic transitions.
Magnetic Properties:
Many transition metal compounds are paramagnetic, meaning they are attracted to a magnetic field.
This is due to the presence of unpaired electrons in the d-orbitals.
High Melting and Boiling Points:
Transition metals generally have high melting and boiling points.
Their strong metallic bonding contributes to this property.
Catalytic Activity:
Transition metals are excellent catalysts for various chemical reactions.
They can facilitate reaction pathways by providing alternative mechanisms with lower activation energy.
4. Coordination Complexes:
Coordination Sphere:
Central metal atom/ion bonded to surrounding ligands.
Ligands are Lewis bases (donate electron pairs).
Coordination Number:
The number of ligands bonded to the central metal.
Common coordination numbers include 2, 4, and 6.
5. Crystal Field Theory (CFT):
Describes the splitting of d-orbitals in a transition metal ion's electron cloud when it is surrounded by ligands.
Two energy levels for d-orbitals: t2g (lower energy) and eg (higher energy).
Splitting is determined by the geometry of the complex (tetrahedral or octahedral).
6. Ligand Field Theory (LFT):
An extension of CFT that takes into account the nature of ligands and their influence on d-orbital splitting.
Strong-field ligands cause larger energy differences between t2g and eg levels than weak-field ligands.
Structure and Shape of Organic Molecules Cheat Sheet
1. Organic Molecules:
Compounds primarily composed of carbon and hydrogen.
Organic chemistry focuses on the structure, properties, and reactions of these compounds.
2. Functional Groups:
Groups of atoms that determine the chemical properties and reactivity of organic compounds.
Examples:
Alkyl (-CH3, -C2H5)
Alkenyl (-CH=CH2)
Alkynyl (-C≡CH)
Hydroxyl (-OH)
Carbonyl (C=O)
Amino (-NH2)
Carboxyl (-COOH)
3. Infrared (IR) Spectroscopy:
Analytical technique used to identify functional groups in organic compounds.
Measures the absorption of infrared radiation by chemical bonds.
Different functional groups have characteristic absorption peaks.
4. Alkanes:
Saturated hydrocarbons with single bonds between carbon atoms.
General formula: CnH2n+2.
Named using IUPAC rules.
Examples: Methane (CH4), Ethane (C2H6), Propane (C3H8).
5. Cycloalkanes:
Saturated hydrocarbons arranged in a closed ring.
General formula: CnH2n.
Examples: Cyclopropane (C3H6), Cyclobutane (C4H8), Cyclopentane (C5H10).
6. Alkenes:
Unsaturated hydrocarbons with at least one carbon-carbon double bond.
General formula: CnH2n.
Named using IUPAC rules.
Examples: Ethene (ethylene, C2H4), Propene (propylene, C3H6), Butene (butylene, C4H8).
7. Chirality:
Property of molecules that are non-superimposable on their mirror images.
Chiral compounds have a "handedness."
Enantiomers are pairs of chiral molecules that are mirror images but cannot be superimposed.
Important in biochemistry and drug development.
8. Molecular Geometry:
Determines the 3D arrangement of atoms in a molecule.
Influences the molecule's properties, reactivity, and polarity.
Common geometries include linear, trigonal planar, tetrahedral, and trigonal bipyramidal.
9. Isomerism:
Compounds with the same molecular formula but different structural arrangements or spatial orientations.
Types:
Structural Isomers: Differ in the connectivity of atoms.
Stereoisomers: Same connectivity but different spatial arrangement (e.g., cis-trans isomers, enantiomers).
10. Nomenclature: - Use IUPAC rules to systematically name organic compounds based on their structure and functional groups. - Prefixes, suffixes, and numbering systems indicate the type and location of substituents.
11. Bonding and Hybridization: - Organic molecules often exhibit sigma (σ) and pi (π) bonds due to the overlap of atomic orbitals. - Hybridization of carbon atoms influences bond angles and shapes.
Atomic Theory Cheat Sheet
1. Dalton's Atomic Theory (1803):
Proposed by John Dalton.
Elements are composed of indivisible particles called atoms.
All atoms of the same element are identical in mass and properties.
Compounds are formed by the combination of atoms in simple, whole-number ratios.
Chemical reactions involve the rearrangement of atoms, and atoms are neither created nor destroyed.
2. Modern Atomic Theory:
Atoms are the basic building blocks of matter.
Atoms are composed of subatomic particles:
Protons (positively charged) in the nucleus.
Neutrons (neutral) in the nucleus.
Electrons (negatively charged) orbiting the nucleus in electron clouds or energy levels.
Elements are defined by the number of protons, known as the atomic number.
3. Subatomic Particles:
Protons:
Mass: ~1 atomic mass unit (amu).
Charge: +1.
Neutrons:
Mass: ~1 amu.
Charge: 0.
Electrons:
Mass: ~1/1836 amu (negligible).
Charge: -1.
4. Atomic Structure:
Atomic Number (Z): The number of protons in an atom.
Mass Number (A): The sum of protons and neutrons in the nucleus.
Isotopes: Atoms of the same element with different numbers of neutrons.
Atomic Mass: Weighted average of the masses of all naturally occurring isotopes of an element.
5. The Periodic Table:
Organizes elements based on atomic number and chemical properties.
Groups (columns) have similar chemical properties.
Periods (rows) represent the number of electron shells.
6. Atomic Models:
Thomson's Model (Plum Pudding Model): Electrons embedded in a positively charged sphere.
Rutherford's Model: Discovered the atomic nucleus, where most of the mass is concentrated.
Bohr's Model: Electrons orbit the nucleus in discrete energy levels or shells.
Modern Quantum Mechanical Model: Describes the behavior of electrons as waves and particles in electron clouds or orbitals.
7. Quantum Numbers:
Principal Quantum Number (n): Represents the energy level or shell (n=1, 2, 3, ...).
Angular Momentum Quantum Number (l): Determines the shape of the orbital (s, p, d, f).
Magnetic Quantum Number (m_l): Specifies the orientation of an orbital in space.
Spin Quantum Number (m_s): Describes the spin of an electron (up or down).
8. Electron Configuration:
The arrangement of electrons in energy levels and orbitals.
Follows the Aufbau principle, Pauli exclusion principle, and Hund's rule.
9. Periodic Trends:
Atomic radius: Increases down a group, decreases across a period.
Ionization energy: Decreases down a group, increases across a period.
Electronegativity: Decreases down a group, increases across a period.
10. Chemical Bonding:
Atoms form bonds to achieve a stable electron configuration.
Types of bonds include covalent, ionic, and metallic.
11. Molecules and Compounds:
Molecules are formed when atoms share electrons (covalent bonding).
Compounds are substances composed of different elements in fixed ratios.
12. Nuclear Reactions:
Involves changes in the nucleus.
Examples include nuclear decay, fusion, and fission.
13. Modern Advances:
The discovery of subatomic particles beyond protons, neutrons, and electrons, such as quarks and neutrinos.
The development of quantum mechanics and its application to atomic theory.
Structure and Bonding Cheat Sheet
1. Types of Chemical Bonds:
Covalent Bonds:
Formed by the sharing of electrons between atoms.
Can be polar (unequal sharing) or nonpolar (equal sharing).
Examples: H2 (hydrogen gas), O2 (oxygen gas).
Ionic Bonds:
Formed when electrons are transferred from one atom to another, resulting in ions.
Ions are held together by electrostatic attractions.
Example: NaCl (sodium chloride).
2. Lewis Dot Structures:
Diagrams representing the valence electrons of atoms.
Useful for predicting and understanding bond formation.
Follows the octet rule: Atoms tend to gain, lose, or share electrons to achieve a full outer shell of 8 electrons (2 for hydrogen).
3. VSEPR Theory (Valence Shell Electron Pair Repulsion):
Predicts molecular geometry based on the arrangement of valence electron pairs around the central atom.
Minimizes repulsion between electron pairs.
Helps determine molecular shape (e.g., linear, bent, tetrahedral).
4. Intermolecular Forces:
Forces of attraction between molecules.
Types:
London Dispersion Forces: Present in all molecules; temporary fluctuations in electron distribution.
Dipole-Dipole Interactions: Occur in polar molecules due to uneven electron distribution.
Hydrogen Bonds: A special type of dipole-dipole interaction involving hydrogen bonded to highly electronegative elements (N, O, F).
5. Valence Bond Theory:
Describes chemical bonding in terms of overlapping atomic orbitals.
Covalent bonds are formed when orbitals of two atoms overlap.
Examples include sigma (σ) and pi (π) bonds.
6. Molecular Orbital Theory:
Describes the behavior of electrons in terms of molecular orbitals (MOs).
MOs are formed by the linear combination of atomic orbitals (LCAO).
Predicts bond order, bond strength, and magnetic properties.
Examples include bonding, antibonding, and nonbonding orbitals.
7. Ionic Compounds:
Composed of cations (positively charged ions) and anions (negatively charged ions).
Held together by electrostatic attractions.
Form crystal lattices.
Conduct electricity when molten or dissolved in water.
8. Covalent Compounds:
Formed by the sharing of electrons.
Can be polar or nonpolar based on electronegativity differences.
Often exist as discrete molecules.
Generally have lower melting and boiling points compared to ionic compounds.
9. Hybridization:
The process by which atomic orbitals mix to form new hybrid orbitals.
Common in organic compounds and explains molecular geometry.
10. Bond Polarity vs. Molecular Polarity:
Bond polarity is determined by the electronegativity difference between bonded atoms.
Molecular polarity depends on the shape and polarity of the bonds within a molecule.
Transition Metals Cheat Sheet
1. Transition Metals:
Located in groups 3 to 12 on the periodic table.
Also known as the "d-block" elements.
Exhibit a wide range of properties due to their partially filled d-orbitals.
Often form colorful compounds and serve as catalysts.
2. Electron Configurations of Transition Metals:
Transition metals have electron configurations that include d-orbitals in addition to s and p orbitals.
The filling order follows the Aufbau principle, Hund's rule, and the Pauli exclusion principle.
For example, the electron configuration of chromium (Cr) is [Ar] 3d^5 4s^1, and not [Ar] 3d^4 4s^2 as might be expected.
3. Properties of Transition Metals:
Variable Oxidation States:
Transition metals can exist in multiple oxidation states due to the availability of d-electrons.
Some elements have particularly stable oxidation states, such as Fe(II) and Fe(III) for iron.
Complex Formation:
Transition metals readily form coordination complexes with ligands (molecules or ions).
Ligands coordinate to the central metal ion through dative covalent bonds.
Complexes often exhibit color due to electronic transitions.
Magnetic Properties:
Many transition metal compounds are paramagnetic, meaning they are attracted to a magnetic field.
This is due to the presence of unpaired electrons in the d-orbitals.
High Melting and Boiling Points:
Transition metals generally have high melting and boiling points.
Their strong metallic bonding contributes to this property.
Catalytic Activity:
Transition metals are excellent catalysts for various chemical reactions.
They can facilitate reaction pathways by providing alternative mechanisms with lower activation energy.
4. Coordination Complexes:
Coordination Sphere:
Central metal atom/ion bonded to surrounding ligands.
Ligands are Lewis bases (donate electron pairs).
Coordination Number:
The number of ligands bonded to the central metal.
Common coordination numbers include 2, 4, and 6.
5. Crystal Field Theory (CFT):
Describes the splitting of d-orbitals in a transition metal ion's electron cloud when it is surrounded by ligands.
Two energy levels for d-orbitals: t2g (lower energy) and eg (higher energy).
Splitting is determined by the geometry of the complex (tetrahedral or octahedral).
6. Ligand Field Theory (LFT):
An extension of CFT that takes into account the nature of ligands and their influence on d-orbital splitting.
Strong-field ligands cause larger energy differences between t2g and eg levels than weak-field ligands.
Structure and Shape of Organic Molecules Cheat Sheet
1. Organic Molecules:
Compounds primarily composed of carbon and hydrogen.
Organic chemistry focuses on the structure, properties, and reactions of these compounds.
2. Functional Groups:
Groups of atoms that determine the chemical properties and reactivity of organic compounds.
Examples:
Alkyl (-CH3, -C2H5)
Alkenyl (-CH=CH2)
Alkynyl (-C≡CH)
Hydroxyl (-OH)
Carbonyl (C=O)
Amino (-NH2)
Carboxyl (-COOH)
3. Infrared (IR) Spectroscopy:
Analytical technique used to identify functional groups in organic compounds.
Measures the absorption of infrared radiation by chemical bonds.
Different functional groups have characteristic absorption peaks.
4. Alkanes:
Saturated hydrocarbons with single bonds between carbon atoms.
General formula: CnH2n+2.
Named using IUPAC rules.
Examples: Methane (CH4), Ethane (C2H6), Propane (C3H8).
5. Cycloalkanes:
Saturated hydrocarbons arranged in a closed ring.
General formula: CnH2n.
Examples: Cyclopropane (C3H6), Cyclobutane (C4H8), Cyclopentane (C5H10).
6. Alkenes:
Unsaturated hydrocarbons with at least one carbon-carbon double bond.
General formula: CnH2n.
Named using IUPAC rules.
Examples: Ethene (ethylene, C2H4), Propene (propylene, C3H6), Butene (butylene, C4H8).
7. Chirality:
Property of molecules that are non-superimposable on their mirror images.
Chiral compounds have a "handedness."
Enantiomers are pairs of chiral molecules that are mirror images but cannot be superimposed.
Important in biochemistry and drug development.
8. Molecular Geometry:
Determines the 3D arrangement of atoms in a molecule.
Influences the molecule's properties, reactivity, and polarity.
Common geometries include linear, trigonal planar, tetrahedral, and trigonal bipyramidal.
9. Isomerism:
Compounds with the same molecular formula but different structural arrangements or spatial orientations.
Types:
Structural Isomers: Differ in the connectivity of atoms.
Stereoisomers: Same connectivity but different spatial arrangement (e.g., cis-trans isomers, enantiomers).
10. Nomenclature: - Use IUPAC rules to systematically name organic compounds based on their structure and functional groups. - Prefixes, suffixes, and numbering systems indicate the type and location of substituents.
11. Bonding and Hybridization: - Organic molecules often exhibit sigma (σ) and pi (π) bonds due to the overlap of atomic orbitals. - Hybridization of carbon atoms influences bond angles and shapes.
