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CHEM 103- Exam 2 Study Guide
Section 5.1: Electromagnetic Radiation
● Electromagnetic radiation occurs daily in real life. We experience electromagnetic radiation when listening to a radio, or using a microwave oven.
● All types of electromagnetic radiation, consists of particles that move as waves of energy. ● The highest part of the wave is known as crest .
● The lowest part of the wave is known as trough.
● Wavelength is the measurement of distance between crest and trough. In other words, wavelength is the distance between the crests and troughs.
● The number of times the crests of a wave pass a point in one second is known as the frequency.
● Frequency is measured in hertz, denoted as “ s ”
● Wave Equation expresses the relationship of the speed of light to wavelength and frequency. You need to memorize this equation for the exam → c = λv
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● The arrangement of different types of radiation from the longest to the shortest wavelength is expressed in electromagnetic spectrum.
● Energy of electromagnetic radiation is directly related to frequency.
● There’s an inverse relationship between the wavelength and frequency: ○ Shorter wavelengths have higher frequencies
○ Longer wavelengths have lower frequencies
○ Different types of electromagnetic radiation have different wavelengths and frequencies We also discuss several other topics like keith pascoe gsu
● Shorter wavelengths are usually produced from infrared lights, ultraviolet light, and heat lamps. Visible light ranges from 400 to 700 nm.
● Radio waves, and waves produced from cellular phones usually have longer wavelengths. ● Blue light has a much more shorter wavelength compared to red light. Blue light is scattered more by the molecules in the atmosphere. Therefore, this is why we we say the sky is blue.
Section 5.2 Atomic Spectra and Energy Levels:
● Continuous Spectrum, is produced when white light from the sun or a lightbulb passes through a prism or raindrops. This is how rainbow, and fireworks are produced. ● Heating atoms of an element produce light
○ Colors are produced when electricity excites electrons in noble gases. ● Photons are stream of particles representing a quantum of light or other electromagnetic radiation.
○ A photon has the characteristics of both a particle and a wave If you want to learn more check out one class notes
○ A photon is a packet of energy that travels at the speed of light as an energy wave
● The energy of a photon is directly proportional to its frequency.
● Photons are important in modern world. They are used in many settings: ○ In hospitals, high energy photons are used in treatments to reach tumors without damaging surrounding normal tissues.
○ Photons are used in lasers to read pits on CDs-DVDs, Don't forget about the age old question of siue change password
● Atomic Spectra is produced when the light emitted from heated elements is passed through a prism.
○ Heated elements produce an atomic spectrum, which consists of different colors separated by dark area.
○ Lines produced in atomic spectra are associated with changes in the energies of electrons.
○ Only certain wavelengths of light are produced by these heated elements ● Flame test experiment demo in class:
○ Strontium (Sr) produced red flame
○ Copper (Cu) produced green flame
○ Sodium (Na) produced orange flame
○ Copper (Cu) produced blue flame
○ Potassium (K) produced violet and purple flame
● In every atom, each electron has its own specific energy level. Energy levels are assigned in values known as the principal quantum numbers
○ As principle quantum numbers increases, this indicates the energy levels are moving farther away from the nucleus.
○ On the other hand, electrons in lower energy levels are closer to the nucleus. ● An electron can have the energy of only one of the energy levels in an atom. The energy of an electron is quantized.
● Changes in energy levels often associate with absorption or emit of energy: ○ When electrons change from a lower to higher energy level, they absorb the energy equal to the change in energy levels.
○ When electrons change from a higher to a lower, they emit energy equal to the change in energy levels.
○ Think of this:
■ Low to high → absorb energy
■ High to low → emit energy
○ High energy indicates higher frequency We also discuss several other topics like itemized deductions are recorded on
● Bohr’s model of the atom shows…
○ Quantized, localized electrons
Section 5.3 Sublevels and Orbitals:
● Energy levels are assigned quantum numbers. As electrons increase in energy, the values of n also increases. Energy levels have a maximum number of electrons equal to 2n .2 “n” represents the energy levels (1-4).
● Quantum Numbers Concept Breakdown:
○ Shell (n) indicates principal level with positive integers (i.e: 1,2,3,4).
○ Subshell (l) indicates angular momentum. (Rule: n-1, where n is shell) ○ Orientation (ml): indicates magnetic number. (Rule: -L, … , +L, where L is subshell)
○ Spin number (ms): indicates spin number. (Rule: +½ OR -½ )
● The maximum number of energy level is 4. The maximum number of electrons that an atom can have is 32.
○ The second energy level, n=2, contains two sublevels: 2s and 2p (I= 0, I= 1). ■ The 2s sublevel just have one orbital with a spherical shape
■ The 2p sublevel contains three p orbitals that are perpendicular to one another.
○ The third energy level, n=3, contains three sublevels: 3s, 3p, and 3d (I= 0, 1, 2) ■ The 3s sublevel just have one orbital with a spherical shape
■ The 3p sublevel have three p orbitals perpendicular to each other → 3px, 3py, and 3pz. If you want to learn more check out rutgers comp apps for business exam 1
■ The 3d sublevel contains five d orbitals
■ There are five different orientations for this sublevel, ranges from (-2 to 2).
○ The fourth energy level, n=4, contains four sublevels: 4s, 4p, 4d, and 4f. ■ The 4s sublevel contains just one orbital with a spherical shape
■ The 4p sublevel contains three p orbitals → 4px, 4py, and 4pz.
■ The 4d sublevel contains five d orbitals
■ The 4f sublevel contains seven f orbitals, with complex shapes (Interval of 3).
● Energy levels consist one or more sublevels, which can be identified by the letters s, p , d , and f. Sublevels contain a number that is equal to the principal quantum number (n). ○ Sublevel s is the lowest in energy, whereas sublevel f is the highest in energy. ○ Order of increasing energy in sublevels: s (lowest) → p → d → f (highest) ● An electron has the highest probability of being found at an orbital, which it’s a three dimensional space around a nucleus. The shape of the orbital represents its electron density. Within any orbital, there could only be 2 electrons found.
● S orbital:
○ For each type of orbital, there will be a specific quantum number. The shape of the orbital is unique to each type of orbital.
○ For instance, in an s orbital (quantum number: 0), electrons will be found in a region with a spherical shape. As the energy level increases, the size of the spherical s orbital increases.
○ For every energy level, there is only one s orbital.
● P orbital:
○ There are total of 3 p orbitals. Beginning with n=2.
○ Each p orbital has two lobes.
○ The p orbitals are arranged perpendicular to each other along the x, y, z axes. The orientation of axes are specified by quantum number with increments of 1.
● According to Pauli exclusion principle, each individual orbital can hold a maximum of two electrons. In other words, no two electrons in a multielectron atom have the same set of four quantum numbers.
● An electron is seen as spinning on its axis and generates a magnetic field. It is specified by another quantum number, for electron spin.
● For each sublevel, there is a limited capacity to which how many number of electrons can occupy in that level.
○ An s sublevel can hold a maximum of one or two electrons.
○ EACH p orbital can hold up to two electrons. Therefore, the three p orbitals in a p sublevel can hold a maximum of six electrons.
○ A d sublevel with five d orbitals can hold a maximum of 10 electrons. ○ A f sublevel with seven f orbitals can hold a maximum of 14 electrons.
● Electrons in the same orbital have opposite spins, and their magnetic fields cancel. To represent the electron spin, we can use arrows.
Section 5.4 Orbital Diagrams and Electron Configurations:
● Orbital diagrams are useful in explaining how electrons are arranged within the atom. Besides that, it also shows:
○ Show the order in which electrons are placed in orbitals
○ Use boxes to represent. Use colors to represent the sublevels s, p, d , and f . ● In order to place electrons in orbital diagrams, we need to…
○ Electrons are represented by arrows, and the direction of the arrow is used to represent electron spin.
○ According to Aufbau principle, electrons need to fill orbitals in order of increasing energy beginning with 1s, then 2s, and 2p.
● Each electron has a unique set of four quantum numbers; an orbital can hold a maximum of two electrons with opposite spins→ Pauli exclusion principle,
● Orbitals can only hold a maximum of two electrons. Moreover, according to Hund’s rule within sublevels that contain multiple orbitals, one electron is placed in each orbital with parallel spins before the electrons are paired because there is less repulsion.
● Electron configuration is a notation widely used by chemists to indicate placement of electrons in an atom.
○ The lowest energy sublevel is written first, then sublevels with increasing energies ○ The number of electrons in each sublevel is written as a superscript
■ Element Hydrogen (H) with atomic number 1
■ Element Helium (He) with atomic number 2
Section 5.5 Electron Configurations and the Periodic Table:
● Different blocks within the periodic table correspond to the s, p, d, and f sublevels. Therefore, this gives up the option of building the electron configurations of atoms by reading the periodic table in order of increasing atomic number.
● Based on the periodic table blocks, we know that s block includes: elements hydrogen and helium, elements in Group 1A, and elements in Group 2A.
○ Elements in the s block have their final one or two electrons in an s orbital. ● P block locates all the way to the right, it consists of elements in group 3A through group 8A.
○ There are six p block elements in each period, this is due to the three p orbitals can hold up to six electrons.
● D block contains the transition elements. There are 10 elements in each d block, because five d orbitals can hold up to 10 electrons.
○ The d sublevels are one less (n-1) than period number
● F block contains the two rows of inner transition elements at the bottom of the periodic table. There are 14 elements in each f block, due to the fact that 7 f orbitals can hold up to 14 electrons.
○ The f sublevel is two less than the periodic number
● Elements with similar properties have the same number of valence electrons. For elements in the “s” and “p” block→ valence electrons are electrons found in the outermost electron shell. This shell has the highest n value.
● The 4s sublevel fills before the 3d sublevel because the electrons in the 4s sublevel are slightly lower than energy than those in the 3d sublevel
○ This order occurs again in Period 5, when the 5s fills before the 4d sublevel, and in period 6, when the 6s fills before the 5d sublevel.
● Some exceptions in sublevel block order
○ In filling the 3d sublevel, exceptions occur for chromium and copper. ○ In chromium, the 3d sublevel is close to being half filled, which is more stable. Thus, one of the 4s electrons moves to the 3d sublevel.
○ In copper, the 3d sublevel is close to being filled, which is more stable. Thus, one of the 4s electrons moves to the 3d sublevel.
Section 5.6 Trends in Periodic Properties:
● Elements that have similar properties have the same of valence electrons. ● For elements in the s and p sublevels → valence electrons are the electrons in the outermost outer electron shell. The shell with the highest energy.
● To determine how many valence electrons an element has, look for the group number on the periodic table
○ Sodium (Na) have one valence electron
○ Aluminum (Al) have three valence electrons
○ Sulfur (S) have six valence electrons (group 6A).
● Valence electrons are depicted in another way known as : Electron dot symbol or Lewis Symbol
○ Electron dot symbol indicates the number of valence electrons as dots around the symbol of the element.
○ For magnesium (Mg), the element has 2 valence electrons, and therefore has two dots on the sides of the symbol Mg. Look at the image to the right, any of these are acceptable electron dot symbol.
● The distance of the valence electrons determine the atomic size of an element ○ Decreases as you move farther to the right of the periodic table. (group number) ○ Increases as you move down from top to bottom of the periodic table. (period number)
● The atomic radius increases….
○ Go down each group of representative elements.
○ The number of energy levels increases
■ As the energy level increases, the valence electrons are farther from the nucleus
● The atomic radius decreases…
○ Go from left to right across the periodic table
○ As more protons increase in the nuclear attraction for valence electrons
● The energy it takes to remove a valence electron. The attraction is to the nucleus → Ionization energy
● When an electron is removed from a neutral atom, a cation with a 1+ charge is formed. ● Ionization energy decreases as the electrons are farther away from the nucleus. In other words, going down a group.
○ Metals: have 1-3 valence electrons. Lower in ionization energies
■ Metals lose their electrons easily. On the other hand, metalloids lose electrons, but not as easily as metals.
○ Nonmetals: have 5-7 valence electrons. Higher in ionization energies ■ Do not easily lose their electrons
○ Noble gases: complete octets. HIGHEST in ionization energies
Practice Question 1: Which element in each pair with the higher ionization energy A. Li or K
B. K or Br
C. P or Cl
Practice Question 2: Place the following in order of decreasing metallic character: Br, Ge, Ca, Ga → Ca, Ga, Ge, Br
Practice Question 3: Going down group 6A (16)
● The ionization energy: Decreases
● The atomic size: Increases
● The metallic character: Increases
● The number of valence electrons: Stays the same
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● Guides to Drawing Orbital Diagrams:
○ Draw boxes to represent the occupied orbitals
○ Place a pair of electrons with opposite spins in each filled orbital
○ Place the remaining electrons in the last occupied sublevel in separate orbitals ● Guides to Writing Electron Configurations:
○ State the number of electrons from the atomic number on the periodic table ○ Write the number of electrons for each orbital in order of increasing energy until filling is complete
○ Write an abbreviated electron configuration by replacing the configuration of the preceding noble gas with its symbol.
● Guides to Writing Electron Configurations Using Sublevel Blocks: ○ Locate the element on the periodic table
○ Write the filled sublevels in order, going across each period
○ Complete the configuration by counting the electrons in the last occupied sublevel block
Section 6.1 Ions: Transfer of Electrons
● Almost all elements other than noble gases are found in combination with other elements in compounds.
● Compounds form when electrons are transferred or shared to give a stable electron configuration to both atoms (a noble gas configuration; octet or duet, a filled valence electron energy level).
● Atoms lose, gain, or share valence electrons in the form of an ionic bond or a covalent bond to obtain an octet (or duet).
● Covalent bonds are formed when atoms of nonmetals share their valence electrons ● Ionic bonds are formed when the valence electrons of a metal are transferred to atoms of nonmetals.
● Ions have electrical charges. Formation of ionic bond requires the loss or gain or valence electrons to form a stable electron configuration (Octet Rule or Duet).
○ Ionization energies of metals in Group 1A-3A are low, therefore metal atoms lose their valence electrons to form positively charged ions.
○ Example 1 : Positive Ions → Loss of Electrons
■ A sodium atom (Na) will lose its 3 valence electrons to form a positively charged sodium ion.
○ Example 2 : Positive Ions → Loss of Electrons
■ A magnesium atom in group 2A has 2 valence electrons. They will obtain a stable electron configuration by losing two valence electrons to form an ion with positive 2+ charge.
○ Example 3: Negative Ions→ Gain of Electrons
■ The ionization energy of a nonmetal atom in Groups 5A(15), 6A(16), and 7A(17) is high.
■ Rather than lose electrons to form ions, a nonmetal atom will gain one or more valence electrons to obtain a stable electron configuration (octet).
○ Example 3: Negative Ions→ Gain of Electrons
■ An atom of chlorine with seven valence electrons gains one electron to form an octet. Because it now has 18 electrons and not 17 electrons, it
becomes a chloride ion (Cl-) with a charge of 1-
○ Positively charged ions of metals are known as cations. Cations generally are smaller than atom because their outermost electrons are removed.
○ Nonmetals with negatively charged ions are called anions. Nonmetal anions are about twice the size of their nonmetal atoms. This is due to extra electrons complete the outermost energy level and cause more repulsion.
○ To determine the charges for each item, we can use group numbers of representative elements in the periodic table to determine their charges.
■ Metal cations (+) charge as their group number (one places if double digit group number)
■ Nonmetal anions (-) charge equal to group number (or -18 for double digit group numbers)
○ Atoms will lose or gain electrons to become like their nearest noble gas atoms.
Section 6.2 Ionic Compounds:
● Ionic compounds are made up of positive and negative ions. These ions are held together by a strong electrical attraction between opposite charges.
● The strong electrical attraction is known as ionic bond.
● The sum of the ionic charges always equal to zero.
● Noble gases do not form ionic compounds because they have a stable electron configuration
● The physical and chemical properties of an ionic compound are very different from those of the original elements
● The chemical formula of a compound gives the symbols and numbers of atoms or ions in the lowest whole number ratio. The sum of the ionic charges is always zero. ● Subscripts in an ionic compound represent the lowest number of positive and negative ions that give an overall charge of zero.
● The group of ions that contains the lowest ratios of ions in an ionic compound is known as the formula unit.
○ Example 1 : Write the formula for the ionic compound, containing the elements Li, and N.
■ Lithium Nitride (Li3N)
○ Example 2: Write the ionic formula for the compound containing Ba and Cl. ■ Ba 2Cl → BaCl2 (Barium Chloride) 2+ −
Section 6.3 Naming and Writing Ionic Compounds:
● Naming of an ionic compound involves two components:
○ The name of the metal ion, and the name of the nonmetal ion
○ Other general rule: a space separates the name of the metal and nonmetal. Subscripts are not used, and they are understood because of the charge balance of the ions in the compound. No prefixes needed!!
● Transitions elements does not follow our rule to figuring out charges of ions based on group numbers.
● The transition elements and representative metals of group 4A (14) and 5A (15), such as Pb, Sn, and Bi can:
○ Lose s electrons as well as d electrons from their highest energy levels ○ Can form two or more positively charged ions, and are described as having a variable charge.
○ Exceptions to the above rules are elements zinc, silver, and cadmium which form only one ion → Zn , Ag , Cd . Zg, Ag, and Cd are transition metals but have 2+ 2+ 2+ a fixed charge so no roman numeral used in name.
○ Sn, Pb, Bi are representative metals in groups 14 and 15 have variable charge, so Roman numeral needed in name.
○ YOU NEED TO memorize exceptions for transition metals, as well as exceptions for representative metals. . However, you do not need to memorize charges because you can use the charge on the nonmetal anions to determine the charge on the metal anion.
● Group of atoms that are covalently bonded with an overall ionic charge is known as a polyatomic ion. You need to memorize these polyatomic ions. The lecturer recommends students to create flashcards to help with memorization.
● Polyatomic ions ending:
○ The ions with more oxygen (ends with -ate)
○ The ions with less oxygen (ends with -ite)
■ “Ate is great, ite is slight”
● Practice Problems:
○ Write the name for each of the following ionic compounds
■ Li2s → Lithium Sulfide
■ SnCl2 → Tin(II) chloride
■ Al2O3→ Aluminum Oxide
■ ZnO → Zinc oxide
Chapter 6.4 Polyatomic Ions:
● A group of atoms that are covalently bonded with an overall ionic charge is known as the polyatomic ion.
○ Many products contain polyatomic ions : fertilizer, and plaster cast.
● Recap from last class regarding naming of polyatomic ions:
○ 1. Polyatomic ions ending in --ate : example: nitrate
○ 2. Polyatomic ions ending in --ite: example: nitrite
○ 3. Polyatomic ions with hydrogen use the prefix hydrogen or bi: example: hydrogen carbonate OR bicarbonate
● To write formulas with polyatomic ions, you need to follow the below steps: ○ Identify the cation and polyatomic ions
○ Balance the charges so that it equals to zero
○ Cation goes first in the formula, using the subscripts from charge balance. ● Example 1: Write the formula for magnesium nitrate
○ Step 1: you need to identify the cation and anion in magnesium nitrate
■ Mg NO
■ Cation Anion
○ Step 2: Then, balance out the charges.
■ Magnesium ion has two positive charges, whereas nitrate ion has 1 negative charge.
■ Therefore, we need to write the formula like this with cations goes first in the formula:
● Mg(NO ) The parentheses enclose the formula of the 3−2
nitrate ion. The subscript outside the parentheses indicates the use
of two nitrate ions to balance the formula.
● (2+ Mg) + 2 (-1 NO) = 0
● To name ionic compounds with polyatomic ions, you need to follow the below steps: ○ First, write the name of the positive ion (usually a metal)
○ Then you write the name of the polyatomic ion last
● Example 1: Name Al (CO ) 2 3 3
○ Answer: Aluminum Carbonate
● Example 2: Name Cu (NO )2 2
○ Step 1: Identify the cation and anion
■ Cu → Cation 2+
■ NO → Anion 2−
○ Step 2: Name out the cation
■ Cu → Copper (II) Ion 2+
○ Step 3: Name out the anion
■ NO → Nitrite Ion 2−
○ Step 4: refer to above step 2-3, together we get the compound name… ■ Copper (II) Nitrite
● Hydrogen First in Compound:
○ If you see hydrogen first in compound, it indicates that when these compounds are dissolved in water, they then will have VERY different properties….They are considered as acids now
○ Keep in mind that rules for naming acids are different for binary acids with just two elements , and ternary acids with three or MORE dif erent elements.
● The naming for Binary Acids, follow the format → Hydro_____ic Acid ○ Binary acids that produce hydrogen ions and a simple nonmetal ion in water named with the prefix hydro and end with --ic acid.
■ HCL → Hydrochloric Acid
● The naming for ternary acid, follow the format → _______ic Acid
○ When the oxygen containing acid from a polyatomic ion ends in --ate, it is named by changing the ate ending to ic acid
● Nitrate ion → Nitric Acid
● Sulfate ion → Sulfuric Acid
● Chlorate ion → Chloric Acid
○ When the oxygen containing acid from a polyatomic ion ends in --ite, it is named by changing the ite ending to ous acid.
● Nitrite Ion → Nitrous Acid
● Sulfite Ion → Sulfurous Acid
● Chlorite Ion → Chlorous Acid
● One exception rule to naming ternary acid… when the ternary acid HCN produces hydrogen ions −
and the simple polyatomic ion CN (ending in --ide). It follows the exact rule for naming binary acids!
○ Example: HCN
■ HCN → Hydrocyanic Acid
Chapter 6.5 Molecular Compounds: Sharing Electrons
● A compound that in which contains two or more nonmetals that form a covalent bond is known as the molecular compound.
○ Valence electrons are shared by nonmetal atoms in a covalent bond to achieve stability (octet or duet).
○ Formation of a molecule occurs when two or more atoms share electrons.
● Formation of the hydrogen H molecule 2
○ Two hydrogen atoms share electrons to form a covalent single bond. ○ Each hydrogen atom acquires two electrons.
○ Each hydrogen atom becomes stable like the nearest noble gas, helium.
● Naming of Molecular Compounds:
○ In molecular compounds, prefixes are used to indicate the number of nonmetals ■ The prefix mono is NOT used on the first nonmetal in the compound. Example: Nitrogen dioxide
Rules to naming molecular compounds:
○ Step 1: name the first nonmetal by its element name
○ Step 2: name the second nonmetal by using the first syllable of its element name followed by ide
○ Step 3: Add prefixes to indicate the number of atoms (subscripts )
Example 1: Name the compound NCl : 3
● Step 1: Identify the nonmetal by its elemental name :
○ In this example, the first nonmetal is nitrogen (N)
● Step 2: Identify the second nonmetal by using the first syllable of its element name followed by ide :
○ The second nonmetal is chloride (Cl).
● Step 3: Add prefixes to indicate the number of atoms (superscripts )
○ One nitrogen : no prefix needed
○ The superscript 3 on the chloride atom is then written as the prefix : tri ● Answer: Nitrogen Trichloride
Example 2: Write the formula for the molecular compound diboron trioxide: ● Step 1: Write the symbols in the order of the elements in the name:
○ B , O
● Step 2: Convert all prefixes to subscripts:
○ B 2O 3
● Answer: B 2O 3
● Guides to writing Ionic Formulas from Ionic Charges:
○ Determine the charges of the metal and nonmetal
○ The metal ion is written first, followed by the ion from the nonmetal. ○ The two Na ions needed for charge balance are represented using a subscript in + the formula
● Guides to Naming Ionic Compounds with Metals That Form a Single Ion ○ Identify the cation and anion
○ Name the cation by its element name
○ Name the anion by using the first syllable of its element name followed by ide ○ Write the name for the cation first and the name for the anion second ● Guides to Naming Compounds with Variable Charge Metals
○ Determine the charge of the cation from the anion
○ Name the cation by its element name and use a Roman numeral in parentheses for the charge
○ Name the anion by using the first syllable of its element name followed by ide ○ Write the name for the cation first and the name for the anion second ● Guides to Writing Formulas from the Name of an Ionic Compound ○ Identify the cation and anion
○ Balance the charges
○ Write the formula, cation first, using subscripts from the charge balance ● Guides to write formulas with polyatomic ions
○ Identify the cation and polyatomic ions
○ Balance the charges so that it equals to zero
○ Cation goes first in the formula, using the subscripts from charge balance. ● Guides to name ionic compounds with polyatomic ions
○ Identify the cation and polyatomic ion (anion)
○ Name the cation using a Roman numeral, if needed
○ Name the polyatomic ion
○ Write the name for the compound, cation first and the polyatomic ion second ● Guides to naming molecular compounds:
○ Name the first nonmetal by its element name
○ Name the second nonmetal by using the first syllable of its element name followed by ide
○ Add prefixes to indicate the number of atoms (subscripts )
● Guides to writing formulas for molecular compounds
○ Write the symbols in the order of the elements in the name
○ Write any prefixes as subscripts
● This chart summarizes chapter 6.
Chapter 10.1 Lewis Structures for Molecules and Polyatomic Ions
● A lewis symbol is a way to represent the valence electrons, which are shown as dots placed on the sides, top, or bottom of the symbol for the element
● A covalent bond forms as Hydrogen atoms move close together to share electrons. Hydrogen is a diatomic molecule.
● A molecule is represented by a lewis structure. The valence electrons of the atoms are arranged to give octets.
● Hydrogen, which has two electrons, only gets a duet
● The shared electrons, or bonding pairs, are shown as two dots or a single line between atoms
● The nonbonding pairs of electrons, or lone pairs, are placed on the outside ● Elements that exist as diatomic molecules → Professor BrINClHOF
● Properties of lewis structure:
○ The sequence of bonded atoms in a molecule or polyatomic ion
○ The bonding pairs of electrons shared between atoms
○ The nonbonding or unshared (lone pairs) of electrons
○ The central atom bonded to other atoms
● Guides to drawing lewis structures:
○ Determine the arrangement of atoms
○ Determine the total number of valence electrons
○ Attach each bonded atom to the central atom with a pair of electrons
○ Place the remaining electrons using single or multiple bonds to complete octets (two for H)
● The electronegativity values of the representative elements in Group 1A (1) to Group 7A (17), which indicate the ability of atoms to attract shared electrons, increase going across a period from left to right, and decrease going down a group.
● The least electronegative element is more willing to share electrons and is found in the center of the lewis structure
● Multiple bonds form when there are not enough valence electrons to complete octets ○ In a single bond, one pair of electrons is shared
○ In a double bond, two pairs of electrons are shared → ex: O2
○ In a triple bond, three pairs of electrons are shared → ex: N2
● Bond length is inversely related to bond strength
○ A double bond is shorter and stronger than a single bond
○ A triple bond is shorter and stronger than a double bond
● Exceptions to the Octet Rule
○ A hydrogen molecule requires just two electrons or a single bond (duet) ○ In BCL3, the B atom has only three valence electrons to share
○ Boron compounds typically have six valence electrons on the central B atoms and form just three bonds
○ P, S, Cl, and Br and I can form molecules in which they share more of their valence electrons. (They do follow octet rule)
■ This expands their valence electrons to 10, 12, or every 14 electrons
■ The P atom in PCl3 has an octet, but in PCl5, the P atom has five bonds with 10 valence electrons
■ In H2S, the S atom has an octet, but in SF6, there are six bonds to sulfur with 12 valence electrons
Chapter 10.2 Resonance Structures:
● Resonance structures consist of two or more lewis structures for the same arrangement of atoms, and are shown with a double headed arrow.
● Keep in mind, resonance structures are written for polyatomic ions with multiple bonds or molecules, and are written by changing the location of a double bond.
● Molecules that contain multiple bonds may contain more than one lewis structure ● Steps to draw resonance structures:
○ Determine the arrangement of atoms. The element with a single atom always go in the middle position.
○ Determine the total number of valence electrons by using the group number on the periodic table.
○ Attach each bonded atom to the central atom with a pair of electrons
○ Lastly, place the remaining electrons as lone pairs around the outside. You need to leave the central atom out till the last. Then use multiple bonds to complete octets if necessary.
Chapter 10.3 Bonding and Properties of Solids and Liquids:
● According to the Valence Shell Electron Pair Repulsion (VSEPR) theory, the three dimensional shape of a molecule is determined by:
○ The number of electron groups surrounding the central atom and the number of atoms bonded to the central atom.
● (VSEPR) Two electron groups
○ Central atoms with two double bonds (two electron groups) have the following properties:
■ Linear electron group geometry
■ Linear shape with a bond angle of 180 degrees to minimize repulsion. ● Repulsion occurs between like charges.
● (VSEPR) Three electron groups
○ Central atoms with two single bonds, one double bond (three electron groups) have the following properties:
■ Trigonal planar electron group geometry
■ Bent shape with a bond angle of 120 degrees to minimize repulsion.
● (VSEPR) Four electron groups
○ Central atoms with four single bonds (four electron groups) have the following properties:
■ Tetrahedral electron group geometry
■ Tetrahedral shape with a bond angle of 109 degrees to minimize repulsion. ○ central atoms with three single bonds, one lone pair (four electron groups) have the following properties:
■ Tetrahedral electron group geometry
■ Trigonal pyramidal shape with a bond angle of 109 degrees to minimize repulsion.
○ Central atoms with two single bonds, two lone pairs (four electron groups) have the following properties:
■ Tetrahedral electron group geometry
■ Bent shape with a bond angle of 109 degrees to minimize repulsion.
● Steps to predicting molecular shape by using VSEPR theory:
○ Draw the electron dot structure (also known as lewis structure)
○ Then, arrange the electron groups around the central atom to minimize repulsion ○ Lastly, use the atoms bonded to the central atom to determine the shape. Chapter 10.4 Bonding and Properties of Solids and Liquids:
● By understanding how electrons are shared in bonds, we can learn more about the chemistry of compounds. Formation of bonds between different atoms also determine how they would share the bonding of electrons.
○ Bonds formed by identical atoms share the bonding electrons equally (nonpolar) ○ Bonds formed between different atoms share the bonding electrons unequally (polar)
● The ability of atoms to attract shared electrons are defined by electronegativity ○ Nonmetals have higher electronegativity.
■ Fluorine has the highest with a value of 4.0
○ Metals have lower electronegativity
■ Cesium and francium have the lowest value of 0.7
● The relative electronegativity value for each element can be predicted by using the periodic table
○ Increases from left to right across a period
○ Decreases from top to bottom on a group
● As mentioned earlier, nonpolar covalent bonds tells us that the bonding electrons are shared equally. On the other hand, polar covalent bonds are bonding electrons that are shared unequally. These concepts can be described by the difference in the electronegativity of the bonding atoms.
○ Nonpolar covalent bond between nonmetals
■ Consists of an equal sharing of electrons
■ Electronegativity difference from 0.0 - 0.4
○ Polar covalent bond between nonmetals:
■ Consists of an unequal sharing of electrons
■ Electronegativity difference from 0.5 - 1.8
● Based on the electronegativity difference, we know that bonds become more polar as the difference in electronegativity increases.
● A polar covalent bond has a separation of charges known as dipole with positive and negative ends represented by the greek letter delta.
● To summarize what we’ve learned in this section, please study the information below: ○ Nonpolar covalent bonds:
■ Electronegativity difference: 0.0 to 0.4
■ Electrons are shared equally
○ Polar covalent bonds:
■ Electronegativity difference: 0.5 to 1.8
■ Electrons are shared unequally
○ Ionic bonds:
■ Electronegativity difference: 1.9 to 3.3
■ The electrons are transferred
Chapter 10.5 Polarity of Molecules:
● All of the bonds in a nonpolar molecule are nonpolar, or the polar bonds/dipoles cancel each other out because of a symmetrical arrangement.
○ Each dipoles have different direction, in which cancels each other out. ○ There is no separation of charge.
● On the other hand, when the dipoles from individual bonds do not cancel each other out, then this is known as the polar molecule.
○ For molecules with two or more electron groups, the shape (such as bent or trigonal pyramidal) determines whether or not the dipoles cancel.
● Steps to determining polarity of a molecule:
○ First, you need to determine if the bonds are polar covalent or nonpolar covalent by referring to the electronegativity difference.
○ Then, if the bonds are polar covalent, then draw the lewis structure and determine if the dipoles cancel out.
○ If the molecule has a bent shape in which the dipoles of the bond do not cancel, then now you can conclude that this molecule is polar.
Section 10.6 Intermolecular Forces between Atoms or Molecules:
● Molecules and ions hold close together in liquids and solids state due to attractive forces. ○ Solids melt and liquids boil when the attractive forces between molecules are broken.
● There are three attractive forces between molecules:
○ Dipole-dipole attractions
■ The positive end of one dipole is attracted to the negative end of a second dipole, such as the attractive forces between two molecules of hydrogen
○ Hydrogen bonding
■ This is the especially strong type of dipole dipole attractions. This
attraction occur between polar molecules that contain hydrogen atoms
bonded to very electronegative atoms (e.g: fluorine, oxygen, and nitrogen) ○ Dispersion forces
■ This is a very weak type of attractive forces that occur between nonpolar molecules. This type of attractive forces occur when movement induces a temporary distortion of the electrons in a molecule, creating a temporary dipole.
■ Dispersion forces make it possible for nonpolar molecules to exist in the form of liquids or solids.
● Learning check 10.6:
○ Indicate type major type of molecule interaction.
■ Nitrogen Trifluoride → Dipole-dipole attractions
■ Chlorine Gas (Cl2) → Dispersion forces
■ HF → Hydrogen bonds
Section 7.1 The Mole
● A mole (mol) is a unit that represents the amount of particles within any element. One mole of any element contains 6.022 × 10 atoms of that element. 23
○ 6.022 × 10 particles represents Avogadro’s number. 23
○ Example: 1 mol of carbon = 6.022 × 10 carbon atoms 23
● To summarize, keep mind that one mole of…
○ An element consists of 6.022 × 10 atoms 23
○ An ion consists 6.022 × 10 ions 23
○ A molecule consists 6.022 × 10 molecules 23
○ A formula unit contains 6.022 × 10 formula units 23
● Avogadro’s number is very important in chemistry because it converts moles of a substance to the number of atoms.
● The subscripts in a formula state the number of moles of each element in one mole of the compound.
● Steps to converting moles/particles of a substance to particles/moles: ○ First, state the given and needed quantities
○ Next, use Avogadro’s number to write conversion factors
○ Last, set up the problem to calculate the number of particles or moles.