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Questions and Answers
Why do longer straight-chain hydrocarbons generally have higher boiling points compared to shorter-chain hydrocarbons?
Why do longer straight-chain hydrocarbons generally have higher boiling points compared to shorter-chain hydrocarbons?
Branched hydrocarbons tend to have lower boiling points than their straight-chain counterparts with similar molecular weights. What is the primary reason for this difference?
Branched hydrocarbons tend to have lower boiling points than their straight-chain counterparts with similar molecular weights. What is the primary reason for this difference?
For a substance to dissolve in a solvent, which of the following must occur?
For a substance to dissolve in a solvent, which of the following must occur?
Why do most ionic compounds dissolve in water?
Why do most ionic compounds dissolve in water?
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Alkanes do not dissolve well in water because:
Alkanes do not dissolve well in water because:
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Which type of solvent would be most effective for dissolving a nonpolar molecule?
Which type of solvent would be most effective for dissolving a nonpolar molecule?
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Which of the following properties is most closely associated with giant ionic structures like sodium chloride?
Which of the following properties is most closely associated with giant ionic structures like sodium chloride?
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Why are metals like copper and sodium good conductors of electricity?
Why are metals like copper and sodium good conductors of electricity?
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Which of the following factors would result in the strongest ionic bond?
Which of the following factors would result in the strongest ionic bond?
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Why do ionic compounds conduct electricity when molten or dissolved in water, but not in the solid state?
Why do ionic compounds conduct electricity when molten or dissolved in water, but not in the solid state?
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When layers of ions slide past each other in an ionic compound, repulsion occurs between ions of like charge. What macroscopic property does this explain?
When layers of ions slide past each other in an ionic compound, repulsion occurs between ions of like charge. What macroscopic property does this explain?
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How does increasing the number of protons affect the ionic radius of isoelectronic ions?
How does increasing the number of protons affect the ionic radius of isoelectronic ions?
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What is the relationship between bond enthalpy and bond length in covalent bonds?
What is the relationship between bond enthalpy and bond length in covalent bonds?
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What is the primary reason for the high melting points observed in ionic compounds?
What is the primary reason for the high melting points observed in ionic compounds?
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In covalent bonding, what determines the bond length between two atoms?
In covalent bonding, what determines the bond length between two atoms?
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How does higher electron density between atoms in a covalent bond affect the bond's characteristics?
How does higher electron density between atoms in a covalent bond affect the bond's characteristics?
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Which molecular geometry results from a central atom with 4 bonding pairs and no lone pairs?
Which molecular geometry results from a central atom with 4 bonding pairs and no lone pairs?
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How do lone pairs of electrons affect the bond angle in a molecule when compared to bonding pairs?
How do lone pairs of electrons affect the bond angle in a molecule when compared to bonding pairs?
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Which of the following best explains why graphite can conduct electricity?
Which of the following best explains why graphite can conduct electricity?
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Why does diamond have such a high melting point?
Why does diamond have such a high melting point?
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Which property of graphene makes it particularly useful in electronics and aircraft?
Which property of graphene makes it particularly useful in electronics and aircraft?
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How does the number of delocalized electrons per atom influence the properties of a metal?
How does the number of delocalized electrons per atom influence the properties of a metal?
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Why are metals typically good conductors of electricity?
Why are metals typically good conductors of electricity?
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What is electronegativity?
What is electronegativity?
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What is a characteristic of a molecule with polar bonds that is also symmetrical?
What is a characteristic of a molecule with polar bonds that is also symmetrical?
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Which of the following represents the correct trend for electronegativity on the periodic table?
Which of the following represents the correct trend for electronegativity on the periodic table?
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Which intermolecular force is present in all molecules, regardless of their structure or polarity?
Which intermolecular force is present in all molecules, regardless of their structure or polarity?
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What leads to stronger London dispersion forces?
What leads to stronger London dispersion forces?
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Which type of intermolecular force is the strongest?
Which type of intermolecular force is the strongest?
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How do intermolecular forces affect the boiling points of simple molecular compounds?
How do intermolecular forces affect the boiling points of simple molecular compounds?
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A substance is known to have hydrogen bonding and London Dispersion forces. What can be said about its boiling point?
A substance is known to have hydrogen bonding and London Dispersion forces. What can be said about its boiling point?
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Flashcards
Ionic Bonding
Ionic Bonding
Electrostatic attraction between oppositely charged ions formed by electron transfer.
Polyatomic Ions
Polyatomic Ions
Ions made up of more than one atom that carry a charge.
Swap and Drop Method
Swap and Drop Method
Technique for determining the formula of ionic compounds by exchanging ionic charges.
Ionic Compounds Conductivity
Ionic Compounds Conductivity
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Brittleness of Ionic Compounds
Brittleness of Ionic Compounds
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Ionic Radius
Ionic Radius
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Covalent Bonding
Covalent Bonding
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Bond Enthalpy
Bond Enthalpy
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London Forces
London Forces
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Branched Hydrocarbons
Branched Hydrocarbons
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Solubility Criteria
Solubility Criteria
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Polar vs Nonpolar Solvents
Polar vs Nonpolar Solvents
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Ionic Compounds in Water
Ionic Compounds in Water
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Hydrogen Bonding
Hydrogen Bonding
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Nonpolar Molecules and Solvents
Nonpolar Molecules and Solvents
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Types of Bonding
Types of Bonding
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Molecular Shape
Molecular Shape
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Bond Angles
Bond Angles
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Lone Pairs
Lone Pairs
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Trigonal Planar
Trigonal Planar
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Tetrahedral
Tetrahedral
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Graphite
Graphite
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Diamond
Diamond
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Metallic Bonding
Metallic Bonding
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Electronegativity
Electronegativity
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Polar Bonds
Polar Bonds
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Boiling Points
Boiling Points
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Seesaw Shape
Seesaw Shape
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Study Notes
Ionic Bonding
- Ionic bonding is the electrostatic attraction between oppositely charged ions, formed when one element loses electrons and another element gains electrons to achieve a full outer shell.
- Group 1 elements form 1+ ions, Group 2 elements form 2+ ions, and so on.
- Group 5, 6, and 7 elements typically gain electrons to form 3-, 2-, and 1- ions, respectively.
- Common polyatomic ions include hydroxide (OH-), nitrate (NO3-), ammonium (NH4+), sulfate (SO42-), and carbonate (CO32-).
- Formula of ionic compounds can be determined using the "swap and drop" method.
- Ionic compounds have giant ionic structures, forming large, regular arrays of alternating positive and negative ions.
- Ionic compounds dissolve in water due to the attraction of polar water molecules to the ions.
- Molten ionic compounds and ionic compounds dissolved in water conduct electricity because the ions are free to move.
- Ionic compounds have high melting points due to the strong electrostatic attractions between ions.
- Ionic compounds are brittle; when struck, layers of ions can slide past each other, causing repulsion between ions of the same charge, resulting in the compound breaking.
- The strength of the ionic bond depends on the size of the charge on the ion and the ionic radius.
- Larger charges and smaller ionic radii result in stronger electrostatic attractions and higher melting points.
- Ionic radius increases down a group due to the addition of electron shells.
- Isoelectronic ions have the same number of electrons but different numbers of protons; as the number of protons increases in isoelectronic ions, the ionic radius decreases due to greater attraction of the electrons to the nucleus.
Covalent Bonding
- Covalent bonding involves the sharing of outer electrons between atoms to achieve a full outer shell.
- Single, double, and triple bonds involve the sharing of two, four, and six electrons, respectively.
- Dative covalent bonds (or coordinate bonds) occur when one atom donates both electrons to another atom or ion.
- Bond enthalpy refers to the strength of a bond, which is inversely proportional to bond length: shorter bonds have higher bond enthalpies.
- Bond length results from a balance between attractive forces between the positively charged nuclei and the shared electrons, and repulsive forces between the nuclei and between the electrons.
- Higher electron density between atoms leads to stronger attractive forces, shorter bonds, and higher bond enthalpies.
Molecular Shapes
- The shape of a molecule is determined by the arrangement of the electron pairs around the central atom, which repel each other equally.
- Lone pairs of electrons repel bonding pairs more strongly than bonding pairs repel each other.
- The number of bonding pairs and lone pairs around a central atom can be determined using dot and cross diagrams.
- The shapes and bond angles of molecules with different numbers of bonding pairs and lone pairs can be summarized in a table:
- No lone pairs:
- 2 Bonding Pairs: Linear (180°)
- 3 Bonding Pairs: Trigonal Planar (120°)
- 4 Bonding Pairs: Tetrahedral (109.5°)
- 5 Bonding Pairs: Trigonal Bipyramidal (120° and 90°)
- 6 Bonding Pairs: Octahedral (90°)
- Lone pairs:
- 3 Bonding Pairs, 1 Lone Pair: Trigonal Pyramidal (107°)
- 2 Bonding Pairs, 2 Lone Pairs: Bent (104.5°)
- 3 Bonding Pairs, 2 Lone Pairs: T-shaped (87.5°)
- 4 Bonding Pairs, 1 Lone Pair: Seesaw (120° and 102°)
- 5 Bonding Pairs, 1 Lone Pair: Square Pyramidal (90° and 81.9°)
- 4 Bonding Pairs, 2 Lone Pairs: Square Planar (90°)
- No lone pairs:
- When working with ions, treat the central atom as having an electron count equal to its group number, adjusted for the ionic charge.
Giant Covalent Structures
- Graphite, diamond, and silicon dioxide are examples of giant covalent structures.
- Graphite has layers of carbon atoms, each bonded to three other carbons, with delocalized electrons between the layers.
- Strong covalent bonds within the layers give graphite a high melting point, while weak forces between layers allow it to slide easily, making it useful in pencils.
- Delocalized electrons make graphite an electrical conductor.
- Diamond has a rigid 3D structure of carbon atoms, each bonded to four other carbons.
- This arrangement gives diamond exceptional hardness, high thermal conductivity, and a high melting point.
- Diamond is an insulator because it lacks delocalized electrons.
- Silicon dioxide has the same structure as diamond and similar properties.
- Graphene is a single layer of graphite, extremely thin, strong, and lightweight, making it useful in electronics, aircraft, and other applications.
Metallic Bonding
- Metallic bonding occurs in metals and involves a giant metallic lattice structure with positive metal ions surrounded by a sea of delocalized electrons.
- The delocalized electrons are attracted to the positive ions, resulting in strong electrostatic attractions.
- The number of delocalized electrons per atom contributes directly to the strength of the metallic bond and the metal's melting point. For example, magnesium (Group 2, donating 2 electrons) has a higher melting point than sodium (Group 1, donating 1 electron).
- Metals are good thermal conductors due to the delocalized electrons transferring kinetic energy.
- Metals are good electrical conductors due to the mobile delocalized electrons carrying charge.
- Metals are malleable and ductile because the positive ions in the lattice can slide past each other while maintaining the electrostatic attractions between the ions and the delocalized electrons.
Electronegativity and Polarity
- Electronegativity is the ability of an atom to attract electrons towards itself in a covalent bond.
- Electronegativity increases up and to the right across the periodic table (excluding noble gases). Fluorine is the most electronegative element.
- The greater the difference in electronegativity between two atoms in a bond, the more ionic character the bond will have.
- If the difference in electronegativity is zero, the bond is purely covalent.
- Polar bonds occur when there is a difference in electronegativity between bonded atoms, causing an uneven distribution of electrons and partial charges.
- Larger differences in electronegativity result in more polar bonds.
- Symmetrical molecules can have polar bonds but be nonpolar overall, such as carbon dioxide, due to the symmetrical distribution of electron density.
- Hydrocarbons, despite having slight electronegativity differences between C and H, are also generally classed as nonpolar.
Intermolecular Forces
- Intermolecular forces are weak attractions that exist between molecules.
- Intermolecular forces are much weaker than the covalent bonds that hold atoms together within a molecule.
- London (or instantaneous dipole-induced dipole) forces:
- Present between all atoms and molecules, even those without permanent dipoles.
- Result from temporary distortions in electron clouds due to the movement of electrons, creating instantaneous dipoles that induce dipoles in neighboring molecules.
- The strength of London forces increases with the size of the molecule or atom, as larger electron clouds are more easily distorted.
- Permanent dipole-dipole forces:
- Exist in molecules with permanent dipoles due to differences in electronegativity.
- These molecules have a partial positive end and a partial negative end.
- Weaker electrostatic attractions occur between the oppositely charged ends of neighboring molecules.
- Molecules also have London forces, but permanent dipole-dipole interactions are stronger.
- Hydrogen bonding:
- The strongest intermolecular force, occurring between molecules containing a hydrogen atom bonded to a highly electronegative atom (N, O, or F).
- Hydrogen bonding occurs due to the attraction between the lone pair of electrons on the electronegative atom and the partially positive hydrogen atom of another molecule.
- Molecules with hydrogen bonding also have London forces and permanent dipole-dipole forces.
Boiling Points and Intermolecular Forces
- Boiling points of simple molecular compounds are largely determined by the strength of intermolecular forces between molecules.
- Stronger intermolecular forces require more energy to overcome, resulting in higher boiling points.
- Longer straight-chain hydrocarbons have more London forces and higher boiling points than shorter-chain hydrocarbons.
- Branched hydrocarbons have weaker London forces due to their inability to pack as closely together, leading to lower boiling points.
Solubility
- Polar substances dissolve well in polar solvents, while nonpolar substances dissolve well in nonpolar solvents.
- For a substance to dissolve, the following must occur:
- Bonds in the solvent must break.
- Bonds in the substance must break.
- New bonds must form between the solvent and the substance.
- The new bonds formed must be at least as strong as the bonds broken for the process to be energetically favorable and for dissolution to occur.
- Polar solvents can be aqueous (containing water which can hydrogen bond) or non-aqueous (containing other polar molecules which can form permanent dipole-dipole interactions).
- Most ionic compounds dissolve in water because the strong attraction of the polar water molecules to the ions overcomes the electrostatic attraction between the ions in the crystal lattice.
- Some non-ionic substances can also dissolve in water, such as alcohols, which can hydrogen bond with water molecules.
- Alkanes, which are nonpolar, do not dissolve well in water, as water forms stronger hydrogen bonds with itself than with the weak London forces of the alkane.
- Nonpolar molecules dissolve best in nonpolar solvents which can form London forces with the solute molecules.
Summary of Bonding
-
The text summarizes the main types of bonding, including:
- Giant covalent structures (graphite, diamond, silicon dioxide)
- Simple molecular structures (ammonia, water)
- Giant ionic structures (sodium chloride, calcium oxide)
- Metallic bonding (magnesium, sodium, copper)
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The text highlights the key properties associated with each type of bonding, including:
- Conductivity (electrical and thermal)
- Solubility
- Melting point
- Boiling point
- Malleability and ductility
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The text emphasizes that polar molecules dissolve best in polar solvents (like water), while nonpolar molecules dissolve best in nonpolar solvents (like hydrocarbons).
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Description
Test your knowledge on ionic bonding concepts and properties. This quiz covers the formation of ions, the characteristics of ionic compounds, and the methods to determine their formulas. Explore the behavior of ionic compounds in different states and their conductivity.
