book Chapter 3 Notes

3. Chemical Compounds

3.1 Classification of Matter

• Most substances are composed of two or more elements combined chemically, called compounds.

  • Examples: hydrogen with oxygen and carbon in water and sugar.

• Learning objectives:

  • Define mixture and compound.

  • Distinguish elements, compounds, and mixtures.

  • Describe element combination to form compounds.

  • Construct systematic names for compounds.

  • Describe characteristics of chemical compounds.

• Review Skills:

  • Particle nature of solids, liquids, and gases.

  • Names and symbols of common elements.

  • Group numbers in the periodic table.

  • Alkali metals, alkaline earth metals, halogens, and noble gases.

  • Metals, nonmetals, and metalloids classification.

  • Nuclear model of the atom.

  • Ions, cations, and anions definitions.

  • Covalent bond, molecule, and diatomic definitions.

  • Covalent bond in a hydrogen molecule, H2.

• Compounds:

  • Substances containing two or more elements.

  • Atoms combine in the same whole-number ratio.

  • Examples: water, sucrose, sodium chloride.

  • Formula for water: $H_2O$$$H_2O$$

  • Formula for sucrose: $$C{12}H{22}O_{11}$$

  • Formula for sodium chloride: $NaCl$$$NaCl$$

• Chemical Formula:

  • Description of a chemical compound's components, using symbols and subscripts for the relative number of atoms.

  • Absence of subscript implies one atom.

• Pure Substances:

  • Elements or compounds with constant composition described by a chemical formula.

  • Example: $$Na2CO3$$ indicates a constant ratio of sodium, carbon, and oxygen atoms (2:1:3).

• Mixtures:

  • Contain two or more pure substances with variable composition.

  • Example: salt water with varying amounts of salt (NaCl) and water ($H_2O$$$H_2O$$).

• Classification of Matter:

  • Pure substance vs. mixture:

    • Constant composition = pure substance.

    • Variable composition = mixture.

    • Describable by a chemical formula = pure substance.

  • Element vs. compound:

    • Describable by a single symbol = element.

    • Chemical formula contains two or more element symbols = compound.

3.2 Compounds and Chemical Bonds

• Chemical bonds:

  • Attractions joining atoms in compounds.

  • Atoms combine in specific ratios, giving constant composition.

• Hydrogen Chloride (HCl):

  • When dissolved in water, it becomes hydrochloric acid.

  • Used in laboratories, food processing, and swimming pool treatment.

• Covalent bond:

  • Sharing of electrons between atoms.

  • Example: bond between hydrogen atoms in $H_2$$$H_2$$ molecules.

• Unequal Sharing of Electrons:

  • In HCl, electrons are shared unequally; more attracted to chlorine.

  • Chlorine gains a partial negative charge ($\delta^−$$$\delta^−$$), hydrogen gains a partial positive charge ($\delta^+$$$\delta^+$$).

• Polar Covalent Bond:

  • Unequal sharing of electrons.

  • Results in partial positive ($\delta^+$$$\delta^+$$) and partial negative ($\delta^−$$$\delta^−$$) charges.

• Nonpolar Covalent Bond:

  • Equal sharing of electrons.

  • Negligible partial charges.

  • Example: bond between hydrogen atoms in $H_2$$$H_2$$.

• Transfer of Electrons:

  • One atom attracts electrons much more strongly than the other.

  • Commonly happens when metallic atoms combine with nonmetallic atoms.

  • Example: Sodium (Na) combines with Chlorine (Cl) to form Sodium Chloride (NaCl).

• Ionic Bond Formation:

  • Sodium atom transfers an electron to a chlorine atom.

  • Sodium loses an electron, becomes a cation with +1 charge: $Na \rightarrow Na^+ + e^−$$$Na \rightarrow Na^+ + e^−$$

  • Chlorine gains an electron, becomes an anion with -1 charge: $Cl + e^− \rightarrow Cl^−$$$Cl + e^− \rightarrow Cl^−$$

  • Attraction between $Na^+$$$Na^+$$ and $Cl^−$$$Cl^−$$ forms an ionic bond.

  • Ionic compounds like NaCl are solids at room temperature, while covalent compounds can be gases or liquids.

• Summary of covalent and ionic bond formation:

  • Electrons shift from one bonding atom to another.

  • Atom attracting electrons more strongly acquires a negative charge.

  • Greater difference in electron-attracting ability = greater electron cloud shift.

  • Sufficiently large difference results in complete electron transfer, forming ions and ionic bonds.

  • Significant, but not complete, transfer = polar covalent bond.

  • Negligible shift = nonpolar covalent bond.

• Predicting bond type:

  • Nonmetal + nonmetal = covalent bond.

  • Metal + nonmetal = usually ionic bond.

3.3 Molecular Compounds

• Molecular Compounds:

  • Composed of molecules

  • Are uncharged collections of atoms held together by covalent bonds

• Ionic Compounds:

  • Contain cations and anions

  • Are held together by ionic bonds

• Molecular vs. Ionic Compounds:

  • All nonmetals in a compound suggests covalent bonds and a molecular compound.

  • Metal-nonmetal combinations suggest ionic bonds and ionic compounds.

• Valence Electrons:

  • Electrons influential in the formation of chemical bonds.

  • Number equals the element’s “A-group” number.

  • Group 7A (F, Cl, Br, I) has seven valence electrons.

  • Group 6A (O, S, Se) has six valence electrons.

  • Group 5A (N, P) has five valence electrons.

  • Group 4A (C) has four valence electrons.

• Electron-Dot Symbol:

  • Visual depiction of valence electrons around an element’s symbol.

  • Also known as electron-dot structures or Lewis electron-dot symbols.

  • Dots placed to the right, left, top, and bottom of the element’s symbol, representing unpaired electrons.

• Octet of Electrons:

  • Having eight valence electrons is stable.

  • Noble gases (group 8A) have an octet (except helium with two electrons).

• Lewis Structure:

  • Elements’ symbols represent atoms, dots represent valence electrons.

  • Way of depicting a molecule.

  • Covalent bonds are usually represented by lines.

• Lone Pairs:

  • Nonbonding pairs of electrons.

  • Each chlorine atom in $Cl_2$$$Cl_2$$ molecule has 1 covalent bond and 3 lone pairs.

• Bonding Preferences:

  • Hydrogen atoms form one covalent bond, achieving two electrons around them.

  • Chlorine atoms form one bond and have three lone pairs.

  • Group 7A elements (F, Br, I) form compounds with one bond and three lone pairs.

  • Group 6A elements (O, S, Se) form two covalent bonds and two lone pairs.

  • Group 5A elements (N, P) form three covalent bonds.

  • Carbon (group 4A) forms four covalent bonds (no lone pairs).

• Hydrocarbons and Organic Compounds:

  • Hydrocarbons contain only carbon and hydrogen.

  • Fossil fuels are primarily hydrocarbons.

  • Organic compounds contain a backbone of carbon-carbon bonds.

  • Organic chemistry is the study of carbon-based compounds.

• Multiple Bonds:

  • Double bonds share four electrons, represented by double lines.

  • Triple bonds share six electrons.

• Drawing Lewis Structures:

  • Give each atom its most common number of covalent bonds and lone pairs.

• Alcohols:

  • Organic compounds with one or more –OH groups attached to a hydrocarbon group.

  • Examples: methanol, ethanol, 2-propanol.

• Molecular Shape:

  • Lewis structures show atom connectivity but not spatial arrangement.

  • Electrons in covalent bonds and lone pairs repel each other.

  • Most stable shape keeps electron groups as far apart as possible.

• Tetrahedral Shape:

  • Four covalent bonds in a methane molecule (CH4) adopt a tetrahedral shape.

  • Bond angles are 109.5°.

• Representations of Methane:

  • Space-filling model: accurate representation of electron-charge clouds.

  • Ball-and-stick model: emphasizes molecular shape and covalent bonds.

  • Geometric sketch: technique for describing tetrahedral structures.

• Ammonia (NH3) and Water (H2O):

  • Lone pairs repel electron groups more strongly than bond pairs.

  • Bond angles in ammonia are about 107°.

  • Water molecule has two covalent bonds and two lone pairs.

  • Angle between bond pairs is about 105°.

• Polarity of Water:

  • Oxygen atoms attract electrons more strongly leading to partial minus charge ($\delta^−$$$\delta^−$$).

  • Hydrogen atoms leading to partial plus charge ($\delta^+$$$\delta^+$$).

  • Attraction between partial positive and negative charges holds water molecules close.

• Structure of Liquid Water:

  • Attractions keep molecules same average distance apart, yet allow constant breaking/forming attractions.

3.4 Naming Binary Covalent Compounds

• Binary Covalent Compounds:

  • Pure substances consisting of two nonmetallic elements.

  • Examples: water ($$H2O$), methane ($$$), methane ($$CH4$$).

• Memorized Names:

  • Some compounds have common names (e.g., water, ammonia, methane) that must be memorized.

  • Organic compounds (methane, ethane, propane) also require memorization initially.

• Systematic Names:

  • Each type of compound has systematic guidelines for naming and formula writing.

  • Binary covalent compounds can be identified by formulas containing just two nonmetallic elements (e.g., $$(N2O3$$)).

• Steps for Naming Binary Covalent Compounds:

  • Prefix for First Element: If subscript > 1, add prefix from Table 3.3 (do not use mono- at beginning).

    • Example: $$N2O3$$ starts with di- (dinitrogen).

  • Name of First Element: Write the name of the first element with the prefix attached.

    • Example: dinitrogen.

  • Prefix for Second Element: Use prefix from Table 3.3, even if subscript = 1.

    • Leave off 'a' or 'o' from prefix before element starting with vowel (oxygen or iodine).

    • Example: dinitrogen tri-

  • Root of Second Element Name: Use root from Table 3.4.

    • Example: dinitrogen triox-

  • Add -ide Ending: Add -ide to the root.

    • Example: dinitrogen trioxide.

  • Hydrogen and Halogens: Compounds with hydrogen and halogens (HX) are often named without prefixes (e.g., hydrogen fluoride for HF).

• Converting Names to Formulas:

  • Reverse the naming steps.

  • Write symbols for elements in order mentioned.

  • Use prefixes to determine subscripts (if no prefix, assume mono-).

3.5 Ionic Compounds

• Ionic compounds:

  • Composed of ions attracted to each other by ionic bonds.

• Cations and Anions:

  • Metallic atoms tend to lose electrons and form cations.

  • Nonmetallic atoms tend to gain electrons and form anions.

  • When a metallic and nonmetallic element combine, electrons transfer to form ions.

  • Positive cations and negative anions attract each other to form ionic bonds.

  • Example: Sodium + Fluorine → Sodium Fluoride (Na+ cation and F− anion).

• Predicting Ionic Charges:

  • Noble gas configuration results in stable electron arrangements.

  • Nonmetallic atoms gain electrons to achieve a noble gas configuration.

    • Halogens (group 17) gain one electron to form -1 ions.

    • Group 16 elements gain two electrons to form -2 ions.

    • Nitrogen and Phosphorus gain three electrons to form -3 ions.

    • Hydrogen gains one electron to form -1 ions.

  • Monatomic anions contain single atoms with a negative charge.

• Cations:

  • Metallic atoms lose electrons to create a cation with the same number of electrons as the nearest noble gas.

    • Alkali metals (group 1) lose one electron to form +1 ions.

    • Alkaline earth metals (group 2) lose two electrons to form +2 ions.

    • Aluminum and group 3 metals lose three electrons to form +3 ions.

  • Monatomic cations are single atoms with a positive charge.

• Naming Monatomic Anions and Cations:

  • Monatomic anions are named by adding -ide to the root of the nonmetal.

  • Names of monatomic cations start with the name of the metal, followed by Roman numeral (charge of the ion).

  • If an element has only one possible charge, the Roman numeral is unnecessary.

    • Group 1 metals, group 2 metals, Aluminum, Zinc, and Cadmium always have fixed charges.

    • Silver usually named as silver ion, not silver(I) ion, as +2 is rare

• Structure:

  • Solid structure of ionic compound is an arrangement of cations and anions.

  • Anions get larger when they gain electron, cations get smaller when they lose electrons

  • Ions arrange to maximize cation-anion attraction and minimize anion-anion and cation-cation repulsions.

• Polyatomic Ions:

  • Charged collection of atoms held together by covalent bonds.

  • Lewis structures are often enclosed in brackets with charge indicated at top right corner.

  • Example: Hydroxide ion ($OH^−$$$OH^−$$).

• Ammonium Ion:

  • Only common polyatomic cation ($NH_4^+$$$NH_4^+$$).

  • Can take the place of a monatomic cation in an ionic crystal structure.

• Formulas to Names:

  • Compounds with symbol for metal and symbol for nonmetal called binary ionic compounds.
    General is MaAb, with M representing a metallic symbol, A nonmetallic and a and b as subscripts
    *Polyatomic Ions:
    Can take the place of monatomic anions, this makes formulas that contain a metallic element and formula for polyatomic ion an ionic compound.

  • Ammonium ion, NH4+, can take place of Matallic cations in an ionic compund, chemical formulas with NH4 with symbol for either nonmetallic or polyatomic represents
    ionic compounds

  • Names consist of name for cation followed by anion.

• Writing formulas:Write the formula, including the charge, for the cation then for the anion. Use subscripts to get an uncharged formula, use lower whole ratio numbers for subscripts, if subscript for polyatomic ion is higher than 1, put formula in parenthesis with subscript outside.