Chemical bonding
60 flashcards covering Chemical bonding for the MCAT Chem / Phys / Psych / Soc section.
Chemical bonding is the process by which atoms join together to form molecules and compounds, creating the building blocks of matter. At its core, it involves the attraction between atoms through the sharing or transfer of electrons. For instance, ionic bonds form when one atom donates electrons to another, resulting in charged particles that attract each other, while covalent bonds involve atoms sharing electrons to achieve stability. This concept is fundamental to understanding chemical reactions, molecular structures, and the properties of substances in everyday life and scientific applications.
On the MCAT, chemical bonding appears frequently in the Chemical and Physical Foundations section, often through multiple-choice questions testing your ability to identify bond types, predict molecular shapes, or explain intermolecular forces like hydrogen bonding. Common traps include confusing polar and nonpolar covalent bonds or overlooking electronegativity differences, so pay close attention to diagrams and calculations. Focus on mastering Lewis structures, VSEPR theory, and hybridization to tackle questions about reactivity and phase changes effectively.
A concrete tip: Practice sketching Lewis diagrams for various molecules to quickly visualize bonding patterns.
Terms (60)
- 01
Ionic Bond
An ionic bond is a chemical bond formed by the electrostatic attraction between oppositely charged ions, typically between a metal and a nonmetal, resulting in the transfer of electrons.
- 02
Covalent Bond
A covalent bond is a chemical bond formed by the sharing of one or more pairs of electrons between two atoms, usually nonmetals, to achieve a stable electron configuration.
- 03
Polar Covalent Bond
A polar covalent bond is a type of covalent bond where electrons are shared unequally between atoms due to a difference in electronegativity, creating partial positive and negative charges.
- 04
Nonpolar Covalent Bond
A nonpolar covalent bond is a covalent bond in which electrons are shared equally between atoms, typically when the atoms have similar electronegativities.
- 05
Metallic Bond
A metallic bond is the attraction between a lattice of positive metal ions and the delocalized electrons that move freely throughout the structure, giving metals their conductivity and malleability.
- 06
Hydrogen Bond
A hydrogen bond is an intermolecular force of attraction between a hydrogen atom bonded to a highly electronegative atom (like oxygen or nitrogen) and another electronegative atom.
- 07
Dipole-Dipole Interaction
Dipole-dipole interactions are attractive forces between the positive end of one polar molecule and the negative end of another polar molecule.
- 08
London Dispersion Forces
London dispersion forces are weak intermolecular attractions caused by temporary dipoles from the uneven distribution of electrons in molecules, present in all substances.
- 09
Van der Waals Forces
Van der Waals forces encompass weak intermolecular attractions, including London dispersion forces and dipole-dipole interactions, that affect the physical properties of substances.
- 10
Electronegativity
Electronegativity is a measure of an atom's ability to attract and hold onto electrons in a chemical bond, influencing the polarity of the bond.
- 11
Lewis Dot Structure
A Lewis dot structure is a diagram that shows the bonding between atoms of a molecule and the lone pairs of electrons, helping to visualize valence electrons.
- 12
Octet Rule
The octet rule states that atoms tend to form bonds until they are surrounded by eight valence electrons, achieving a stable electron configuration like noble gases.
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VSEPR Theory
VSEPR theory predicts the shape of molecules based on the idea that electron pairs around a central atom repel each other and arrange to minimize repulsion.
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Molecular Geometry
Molecular geometry describes the three-dimensional arrangement of atoms in a molecule, determined by the positions of bonding and lone pairs around the central atom.
- 15
Bond Angle
A bond angle is the angle formed between two adjacent bonds in a molecule, influenced by the repulsion between electron pairs as per VSEPR theory.
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Hybridization
Hybridization is the concept where atomic orbitals mix to form new hybrid orbitals that explain the geometry and bonding in molecules, such as sp, sp2, or sp3.
- 17
Sigma Bond
A sigma bond is the strongest type of covalent bond, formed by the head-on overlap of atomic orbitals, allowing for free rotation around the bond.
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Pi Bond
A pi bond is a covalent bond formed by the sideways overlap of p orbitals, resulting in electron density above and below the bond axis, typically found in double and triple bonds.
- 19
Resonance Structures
Resonance structures are different Lewis structures for a molecule that differ only in the placement of electrons, not atoms, and represent the delocalized nature of electrons.
- 20
Formal Charge
Formal charge is a method to determine the most stable Lewis structure by calculating the charge on an atom as valence electrons minus nonbonding electrons minus half of bonding electrons.
- 21
Bond Dissociation Energy
Bond dissociation energy is the energy required to break a specific bond in a molecule, indicating the strength of that bond.
- 22
Bond Enthalpy
Bond enthalpy is the average energy needed to break a particular type of bond in a molecule, used to calculate the enthalpy change in reactions.
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Lattice Energy
Lattice energy is the energy released when gaseous ions form a solid ionic compound, reflecting the strength of ionic bonds in the crystal lattice.
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Electronegativity Difference
The electronegativity difference between two atoms determines the bond type: less than 0.4 is nonpolar covalent, 0.4 to 1.7 is polar covalent, and greater than 1.7 is ionic.
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Pauling Scale
The Pauling scale is a numerical scale that assigns electronegativity values to elements based on bond dissociation energies, with fluorine having the highest value.
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Percent Ionic Character
Percent ionic character estimates how ionic a bond is based on electronegativity differences, with higher percentages indicating more ionic behavior.
- 27
Dipole Moment
A dipole moment measures the polarity of a molecule, calculated as the product of charge separation and distance, and is a vector quantity.
- 28
Coordinate Covalent Bond
A coordinate covalent bond is a covalent bond in which one atom donates both electrons to the shared pair, often seen in complex ions.
- 29
Network Covalent Solids
Network covalent solids are substances where atoms are bonded covalently in a continuous network, like diamond, resulting in high melting points.
- 30
Molecular Solids
Molecular solids consist of molecules held together by intermolecular forces, leading to lower melting points compared to ionic or covalent network solids.
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Ionic Solids
Ionic solids are crystalline structures of ions held together by ionic bonds, which conduct electricity when melted or dissolved.
- 32
Metallic Solids
Metallic solids are composed of metal atoms with metallic bonds, allowing for high electrical and thermal conductivity.
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Exceptions to Octet Rule
Exceptions to the octet rule occur in molecules like BF3 or SF6, where atoms have fewer or more than eight valence electrons due to electron deficiency or expansion.
- 34
Hypervalency
Hypervalency refers to elements, typically from period 3 and beyond, that can have more than eight valence electrons by using d orbitals for bonding.
- 35
Bond Order
Bond order is the number of chemical bonds between a pair of atoms, calculated in molecular orbital theory as half the difference between bonding and antibonding electrons.
- 36
Molecular Orbital Theory
Molecular orbital theory describes bonding by combining atomic orbitals to form molecular orbitals, which can be bonding, antibonding, or nonbonding.
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HOMO
HOMO, or highest occupied molecular orbital, is the molecular orbital with the highest energy that contains electrons, influencing chemical reactivity.
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LUMO
LUMO, or lowest unoccupied molecular orbital, is the molecular orbital with the lowest energy that does not contain electrons, important for reactions.
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Antibonding Orbital
An antibonding orbital is a molecular orbital higher in energy than the atomic orbitals from which it forms, which weakens the bond when occupied.
- 40
Bonding Orbital
A bonding orbital is a molecular orbital lower in energy than the atomic orbitals, stabilizing the molecule by holding atoms together.
- 41
Electron Affinity
Electron affinity is the energy change when an atom gains an electron, affecting how atoms form bonds, especially in ionic compounds.
- 42
Ionization Energy
Ionization energy is the energy required to remove an electron from an atom, influencing the formation of ions in chemical bonding.
- 43
Common Trap: Ionic vs. Covalent Bonds
A common mistake is confusing ionic bonds, which involve electron transfer between metals and nonmetals, with covalent bonds, which involve electron sharing.
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Strategy for Drawing Lewis Structures
To draw Lewis structures, first count valence electrons, arrange atoms, distribute electrons to satisfy the octet rule, and add multiple bonds if needed.
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Worked Example: Lewis Structure of Water
For water (H2O), oxygen has six valence electrons and each hydrogen has one; the structure shows oxygen with two single bonds to hydrogens and two lone pairs.
- 46
Worked Example: VSEPR for Ammonia
In ammonia (NH3), the central nitrogen has three bonding pairs and one lone pair, resulting in a trigonal pyramidal geometry with bond angles of about 107 degrees.
- 47
Boiling Point Trends and Intermolecular Forces
Boiling points increase with stronger intermolecular forces; for example, hydrogen bonding in water leads to a higher boiling point than in similar nonpolar molecules.
- 48
Hydrogen Bonding in DNA
Hydrogen bonding between base pairs in DNA, such as between adenine and thymine, stabilizes the double helix structure of the molecule.
- 49
Bond Length
Bond length is the distance between the nuclei of two bonded atoms, shorter for stronger bonds and influenced by atomic size and bond type.
- 50
Bond Strength
Bond strength refers to the energy required to break a bond, with triple bonds being stronger than double bonds, which are stronger than single bonds.
- 51
Fajans' Rules
Fajans' rules predict whether a bond will be more ionic or covalent based on ion size and charge, with smaller, highly charged ions favoring covalent character.
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Common Trap: Resonance vs. Isomers
Resonance structures are different electron arrangements in the same molecule, unlike isomers which are distinct compounds with the same formula.
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Strategy: Determining Molecular Polarity
To determine if a molecule is polar, check for polar bonds and asymmetrical geometry; if both are present, the molecule has a net dipole moment.
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Worked Example: Formal Charge Calculation
For the nitrate ion (NO3-), formal charge on nitrogen is calculated as 5 - 0 - (1/2)8 = +1, helping to identify the most stable resonance structure.
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Allotropes and Bonding
Allotropes like diamond and graphite differ in bonding; diamond has a tetrahedral network of covalent bonds, while graphite has layered hexagonal sheets.
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Intermolecular vs. Intramolecular Forces
Intermolecular forces act between molecules, like hydrogen bonds, while intramolecular forces are the bonds within a molecule, such as covalent bonds.
- 57
Polarity of Molecules
The polarity of a molecule depends on both the presence of polar bonds and the molecule's shape; symmetrical molecules with polar bonds can be nonpolar overall.
- 58
Sigma and Pi Bonds in Ethene
In ethene (C2H4), the double bond consists of one sigma bond from sp2 orbital overlap and one pi bond from p orbital overlap.
- 59
Hybridization in Methane
Methane (CH4) exhibits sp3 hybridization, where carbon's s and three p orbitals mix to form four equivalent orbitals for tetrahedral bonding.
- 60
Resonance in Benzene
Benzene has two resonance structures with alternating double bonds, representing a delocalized pi electron system that stabilizes the ring.