Chem ch. review (6, 12, 13) Cracolice
carboxyl group
(COOH) is a functional group consisting of a carbonyl group (C=O) with a hydroxyl group (O-H) attached to the same carbon atom.
single bond
-Bonds between atoms can be formed by the sharing of one, two or three electron pairs. - composed of one electron pair shared between two atoms and is often represented as a pair of dots or a slash / line.
linear shape
-any two atoms (ex: O2, H2, Br2) -linear w/3 elements (no unshared e-) (ex. CO2) -no unshared pairs, double bonds, one central atom -180 degrees
covalent bond
-atoms share one or more pairs of electrons with other atoms. -degree of sharing varies -even sharing where maximum electron density is located half way between the two atoms -uneven sharing where the maximum electron density is located nearer to one atom than the other. -The simplest covalent bond is formed between a pair of hydrogen atoms. The electron cloud becomes concentrated between the two hydrogen Gilbert N. Lewis proposed that two atoms in a molecule are held together by a covalent bond.
ionic bonds
-between metals and non metals -opposites attract (electrostatic attraction) -transfer of one or more valence electrons from one atom to another (to the nonmetal becoming anion/cation) -ion arrangement is geometric (crystals) All ionic compounds are made from ionic bonds. • Positive cations and the negative anions are attracted to one another through electrostatic forces. • Opposites Attract!
ionic compounds
-solid ionic compounds are insulators -have high melt points due to strong bond forces -ions in rigid locations -metals form positive charges, nonmetals form negative All ionic compounds are made from ionic bonds. usually formed between metals and nonmetals (opposite ends of the periodic table). • If ionic compounds are heated high enough to melt or are dissolved into water, they become conductors of electricity.
drawing lewis diagrams
1. Sum all the valence electrons for each atom in the molecule. The number of valence electrons equals the group number. 2. Connect the central atoms to the peripheral atoms by single bonds. The central atom is the one present in the lowest amount. Indicate single bonds as a line or dash. 3. Complete the octet on the peripheral atoms except hydrogen which is complete with two electrons in its outer shell. An octet has eight electron around the atom. 4. Place any left over electrons on the central atoms. 5. If the central atom lacks an octet, try multiple bonds by converting a lone pair on a peripheral atom into a bond pair.
naming acids
1. Write down the stem of the anion. 2. Add the suffix (-ic) to the anion stem. 3. Add the prefix (hydro-) to the anion stem. 4. Add the word acid to complete the name. ex: HCl Chlor Chloric Hydrochloric Hydrochloric acid ex: H2S Sulf Sulfuric Hydrosulfuric Hydrosulfuric acid
writing ionic compounds using Lewis dot
1. Write the Lewis dot structures of each element 2. Determine the number of valence electrons around each element 3. If element has 3 or less electrons, it has to lose all the electrons to achieve octet 4. If element has 5 or more electrons it has to gain all electrons needed to achieve octet. 5. Determine number of electrons that need to be lost or gained by the elements. 6. Transfer the necessary number of electrons. 7. If octet is not achieved, use multiple numbers of each element until octet is achieved. 8. Write formula showing ratio of each ion needed to make a neutral compound. Cation written first, followed by anion.
writing electron configurations
1. write the number of the shell 2. write the letter of the orbital 3. write the number of electrons in that orbital
tetrahedral
109.5 degrees (each bond has this angle) 4 bonds to a central atom usually in group 4 and 14 elements
electron orbital configurations
1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s 7p to figure out the order of the configuration, draw an arrow on the left diagonal (arrow points left) ex: 1s 2s 2p 3s 3p 4s 3d 4p 5s
number of electron pairs and bond angle
2 electron pairs = 180 degrees (linear) 3 electron pairs = 120 degrees (triangular, trigonal planar) 4 electron pairs = 109.5 degrees (tetrahedral)
Electron pair geometry 3 electron groups
3 electron pairs around center/atom 3 bonded pairs 0 lone pairs triangular planar - e-pair domain geometry triangular planar - molecular geometry 120 degrees ex. BH3 ------------- 3 electron pairs around center/atom 2 bonded pairs 1 lone pair triangular planar - e-pair domain geometry bent - molecular geometry 120 degrees ex. SO2
electron pair geometry 4 electron groups
4 electron pairs around center/atom 4 bonded pairs 0 lone pairs tetrahedral - e-pair domain geometry tetrahedral - molecular geometry 109.5 degrees ex. CH4 ----------- 4 electron pairs around center/atom 3 bonded pairs 1 lone pair tetrahedral - e pair domain geometry trigonal/triangular pyramidal - molecular geometry 107 degrees ex. NH3 ----------- 4 electron pairs around center/atom 2 bonded pairs 2 lone pairs tetrahedral - e pair domain geometry bent - molecular geometry 104.5 degrees ex. H20
monatomic gases
8A: He, Ne, Ar, Kr, Xe, Rn
anhydrous compound
A compound with all water removed, especially water of hydration. For example, strongly heating copper(II) sulfate pentahydrate (CuSO.
polarity of molecules
A homonuclear diatomic molecule in which the bond pair is shared equally between the two atoms has a symmetrical electron distribution and is nonpolar. In a diatomic molecule, where the atoms have different electronegativities the bond pair is held closer to the more electronegative atom. The polarity of polyatomic molecules is determined by a combination of bond polarity and molecular geometry. If a molecule has polar bonds and a geometry such that the charge on the central atom is negative, the molecule is polar. H2O and NH3 are examples. If a molecule has polar bonds and a geometry such that the center is a positive charge and the intersection of negative charge are in the same location, the molecule is nonpolar.
Monatomic Ions anion
A negatively charged ion. • Formedwhenaneutralatomgainsanelectron: There are more electrons (-) than protons (+). F+e- F- 1s22s22p5 1s22s22p6 Electron configuration of Neon : Ne 1s22s22p6 • F- has the same electronic configuration and number of electrons as Ne
ammonium ion
A positively charged ion, NH4, derived from ammonia and found in a wide variety of organic and inorganic compounds. Compounds of ammonium chemically resemble the alkali metals.
monoatomic ion cation
A positively charged ion. Formed when a neutral metal atom loses an electron. There are more protons (+) than electrons (-). Example: Na Na+ + e- 1s2 2s2 2p6 3s1 1s2 2s2 2p6 Electron configuration of Neon Ne 1s22s22p6 Na+ has the same electronic configuration and number of electrons as Ne
Condensed structural formulas
A structural representation of a compound that includes all of the atoms present in a molecule or other chemical entity but represents only certain bonds as lines in order to emphasize a structural characteristic.
elements existing as triatomic or higher molecules
A triatomic molecule is a molecule consisting of 3 atoms (the same or different). Examples of triatomic molecules include (but not limited to): Water (H2O) - Ozone (O3) - Carbon dioxide (CO2)
molecules with more that four electron pairs around the central atom
Atoms of elements in the third period and higher can have more than four electron pairs surrounding them. This is accomplished by involving d-orbitals in the formation of covalent bonds. Phosphorus pentafluoride, PF5 places five electron-pair bonds around the phosphorus atom; six pairs surround sulfur in SF6.
incomplete octet elements
Boron Beryllium
trigonal planar
Boron is the exception to the octet rule usually seen in group 3 and 13 elements no unshared electron pairs bond angle is 120 degrees apart from each other ex. BF3
hydrates
Certain ionic compounds when formed by crystallization from aqueous solutions incorporate a definite number of molecules of water into their formulas. These molecules of water are included in the writing of the formulas. Examples of such hydrates include CuSO4 • 5H2O, CoCl2 • 6H2O, and FeSO4 • 7H2O. If these hydrates are heated, they lose their waters of hydration and may exhibit a change in color.
double bond
Double Bonds - are composed of two electron pairs shared between two atoms and are represented as four dots or as a pair of lines.
procedure for drawing Lewis diagrams - Find the total number of electrons: Tabulate the total number of outer energy level electrons for all atoms in the molecule. For each atom, read the group number.
Draw a first tentative structure: The element with the least number of atoms is usually the central element. Draw a tentative molecular and electron arrangement attaching other atoms with single bonds as the first guess. Single bonds represented with a line represent 2 electrons -Add electrons as dots to get octets around atoms: When counting electrons for the octet around an atom, count both electrons in a bond for each atom and any lone pair electrons. Hydrogen, of course, gets only 2 electrons. Count the total number of electrons in the final structure to see if the total agrees with the number tabulated in step #1. If not, then move a lone pair of electrons into a double bond. Or add more lone pairs of electrons. -Cycle through steps 3 and 4 several times until you get it right by trial and error.
Four electron pairs
Electron pair geometry is tetrahedral. Bond angles are tetrahedral (109.5 degrees). (i) all electron pairs bonding, molecular geometry is tetrahedral. (ii) three electron pairs bonding, one lone pair: molecular geometry is trigonal pyramidal. (iii) Two electron pairs bonding, two lone pairs, molecular geometry is bent (angular).
chemical family
Elements in the same _________ form monoatomic ions having the same charge. They do this by gaining or loosing electrons until they become isoelectronic with the noble gas.
diatomic molecules
For substances that commonly exist as diatomic molecules, it is necessary to distinguish between atoms and molecules and to use the appropriate symbol to refer to atoms of the element or formula to refer to an actual sample of the element under normal conditions.
Acids whose Anions are Oxyanions
For the oxyanion with the greater number of oxygen atoms: a) Name the oxyanion. b) Replace the -ate from the oxy anion name with -ic acid 2. For the anion with the lesser number of oxygen atoms: a) Name the oxyanion. b) Replace the -ite from the oxy anion name with -ous acid Example : HNO3 (a) nitrate (b) nitric acid Example : HNO2 (a) nitrite (b) nitrous acid Example : SO4 (a) sulfate (b) sulfuric acid Example : SO3 (a) sulfite (b) sulfurous acid
trends in electronegativity
From left to right across a period of elements, electronegativity increases. If the valence shell of an atom is less than half full, it requires less energy to lose an electron than to gain one. Conversely, if the valence shell is more than half full, it is easier to pull an electron into the valence shell than to donate one. From top to bottom down a group, electronegativity decreases. This is because atomic number increases down a group, and thus there is an increased distance between the valence electrons and nucleus, or a greater atomic radius. Important exceptions of the above rules include the noble gases, lanthanides, and actinides. The noble gases possess a complete valence shell and do not usually attract electrons. The lanthanides and actinides possess more complicated chemistry that does not generally follow any trends. Therefore, noble gases, lanthanides, and actinides do not have electronegativity values. As for the transition metals, although they have electronegativity values, there is little variance among them across the period and up and down a group. This is because their metallic properties affect their ability to attract electrons as easily as the other elements.
Number of Hydrogen atoms in an Acid
General form of Acid HX- Example: HCl • monoprotic acid • contains one ionizable hydrogen atom General form of Acid H2X - Example: H2SO4 • diprotic acid • contains two ionizable hydrogen atoms General form of Acid H3X -Example: H3PO4 • triprotic acid • contains three ionizable hydrogen
following elements normally exist as diatomic molecules.
H2 N2 O2 F2 Cl2 Br2 I2
noble gases
He Ne Ar Kr Xe Rn
polar covalent bond
If the atoms in the bond pair are not identical, then the bonding electrons are displaced towards the atom that has the greater affinity for electrons and results in a charge displacement. The electrons are pulled closer to the nucleus with greater attracting strength.
naming
If the subscript of the first element is 2 or more, you write the appropriate prefix We have 1 Carbon atom therefore is no prefix Then write the name of the first element in the formula Carbon Write the prefix that denotes the number of atoms of the second element prefix- di, then write the "stem" of the second element in the formula and a suffix (-ide). oxide Name : Carbon dioxide CO2
nonpolar covalent bond
In a diatomic molecule composed of two atoms of the same element, the bonding electrons are shared equally. The nucleus of each atoms is equally attracting the shared electrons.
force that holds atoms together in a covalent bond
In covalent bonding, both atoms are trying to attract electrons--the same electrons. The electrons are shared tightly between the atoms. The force of attraction that each atom exerts on the shared electrons is what holds the two atoms together.
naming non metals
In order to name molecules (made up of non metals) : 1. If the subscript of the first element is 2 or more, you write the appropriate prefix (prefix- di, tri, tetra etc) 2. Then write the name of the first element in the formula 3. Write the prefix that denotes the number of atoms of the second element (prefix- mono, di, tri, tetra etc) 4. Then write the "stem" of the second element in the formula and a suffix (-ide).
naming ionic compounds
In order to name molecules (made up of non metals) : 1. Name the cation first 2. Then write the stem name of the anion in the formula and add the suffix (-ide). 3. Example K2O Cation : Potassium Anion:OxygenStem:Ox + ide!Oxide Name of ionic compound : Potassium Oxide
writing ionic compounds using periodic table
Ionic compounds can be written by considering the charges on ions within the compound 1. Identify the element that becomes the cation 2. Determine the charge on the element by reading group number 3. Identify the element that becomes the anion 4. Determine the charge on the element by reading group number and subtracting 8 from the group number. 5. Write the cation first, followed by the anion 6. Switch the charges between the ions and write the charges as subscripts after the symbol of the element. 7. Omit the positive and negative sign from the charges.
naming oxy anions
Knowing how to name the oxyanions is important to naming the acid containing the oxyanion. ●oxyanion containing the greater number of oxygen atoms has the ending (-ate) ●oxyanion containing the greater number of oxygen atoms has the ending (-ite) ex: NO3- (nitrate) NO2- (nitrite)
Lewis diagrams
Molecules can be depicted by Lewis Diagrams by placing dots or lines around the constituent elemental symbols. Once again only valence electrons are shown. Lines denote bonded electron pairs, whereas dots are reserved for unbounded electrons. The following algorithm can be used to construct Lewis diagrams of most molecules. 1. Find the total number of electrons: Determine the total number of valence electrons by reading the group number for each element. 2. Draw a first tentative structure: Frequently the central element is the one with least atoms in the molecule. In your first draft, attach the atoms with single bonds. 3. Add electrons as dots to get octets around atoms: Each atom must have its valence shell completely filled. Remember to count bonds and lone electrons. 4. Count the total number of electrons: Ensure the number of electrons displayed in the symbol agree with the calculation from step 1. 5. Cycle through steps 3 and 4: By way of trial and error, repeat the steps above until a conclusion is reached.
ions that don't conform to rules
NH4+ ammonium ion OH- hydroxide ion C2H3O2- acetate ion CN- cyanide ion CrO42- chromate ion Cr2O72- dichromate ion MnO4- permanganate ion
common stems
Name of element and its stem Chlorine - chlor Iodine - Iod Sulfur - sulf Fluorine -fluor Nitrogen -nit Selenium - selen Oxygen - ox Phosphorus - phosph Carbon - carb Bromine - brom
naming of monoatomic anions
Naming of Monatomic Anions - are named by adding the ending (-ide) to the elemental stem. ex: Group 7A non metals Chlorine to Cl- to Chloride (anion name) Oxygen to O2- to oxide (anion name) Nitrogen to N3- to nitride (anion name) Sulfur to S2- to sulfide (anion name)
alcohol
O-H) any organic compound whose molecule contains an alkyl group and one or more hydroxyl groups attached to a carbon atom (general formula R
hydroxyl group - an organic functional group,
OH, characteristic of alcohols.
acids and anions derived from total ionization
Some Acids contain polyatomic ions. If the polyatomic ions contain oxygen, they are called oxyanions. Acids that contain oxyanions are named differently as opposed to acids whose anions don't contain oxygen.
odd electron molecule
The bonds in which odd numbers of electrons are used in its formation are known as odd-electron bonds and the molecules are called odd-electron molecules. These molecules are quite stable and paramagnetic in nature. Molecules with an odd number of electrons disobey the octet rule.
Electron pairs (groups)
The electron pair geometry (arrangement of electron pairs) and eventually the molecular geometry (shape) of molecules are determined by counting the number of electron pair groups around a selected central atom and matching them with the assigned geometry. Counting electron pair groups A single bond is counted as 1 electron group A double bond is counted as 1 electron group A triple bond is counted as 1 electron group A lone pair electron is counted as 1 electron group
nonbinding/lone pairs
The electron pairs isolated to a single atom and not shared in bonding
use of the periodic table to predict charge on a monatomic ion
The electronic configuration of many ions is that of the closest noble gas to them in the periodic table.
s and p subshells
The elements in the s and p subshells (group IA-VIIA) tend to lose, gain or share electrons in order to achieve an octet in their outermost shell. • Elements that lose electrons now have an excess of positive charges and form ions called cations. • Elements that gain electrons now have an excess of negative charges and form ions called anions.
hydroxide ion
The hydroxide ion is a negatively charged molecule made up of one oxygen bonded to one hydrogen. When dissolved in water, the hydroxide ion is an incredibly strong base. In fact, according to the Arrhenius definition of a base, the presence of a hydroxide ion is what makes a chemical a base.
naming of monatomic anions
The monatomic anions are named by adding ide to the root of the name of the nonmetal that forms the anion. For example, N3- is the nitride ion.
naming of hydrates
The name of a hydrate follows a set pattern: the name of the ionic compound followed by a numerical prefix and the suffix -hydrate
resonance hybrid
The net sum of valid resonance structures is defined as a resonance hybrid, which represents the overall delocalization of electrons within the molecule. A molecule that has several resonance structures is more stable than one with fewer.
Monatomic Ions with Noble Gas Electron Configurations
The number of electrons lost or gained depends on an element's position in the periodic table relative to the nearest inert gas. • Elements in group IA, IIA and IIIA tend to lose electrons to have the electronic configuration of a noble gas element. • Elements in group VA, VIA and VIIA tend to gain electrons to have the electronic configuration of a noble gas element.
Molecular Geometry/ Shape of molecule
The overall shape of a molecule is defined by: -the bond lengths - the bond angles - the presence or absence of lone pair electrons • The molecular geometry that also describes the shape only takes into consideration the bonded electron pair groups around the central atom. • The lone pair electrons affect the bond angles, but they are not counted when describing the shape or molecular geometry around the central atom. • If there are no lone pair electrons, then the molecular geometry is the same as the electron pair domain geometry. • The presence of lone pair electrons exert a compressive force on bond angles causing them to be smaller than ideal. • The absence of an atom in a position occupied by a lone pair also alters the shape or geometry of the molecule because shape is determined by connecting the points together if one of the points is missing, then the overall shape is alter.
bond/electron pair
The slash of symbol (-) between atoms are referred to as bonding pairs of electrons. • 1 bond = 2 electrons
difference in electronegativity
The smaller the difference in electronegativity, the more non polar the bond is. Electrons are shared more equally. The larger the difference in electronegativity the more polar the bond Electrons are not shared equally. Absolute difference in electronegativity 0.0-0.4 = non polar covalent 0.5-1.9 = polar covalent 2.0-3.3 = ionic
basic idea of VSEPR
Used to predict the geometry of individual molecules from the number of electron pairs surrounding their central atoms.
meaning of acronym VSEPR
Valence shell electron pair repulsion (VSEPR) theory is a model used, in chemistry.
hydrocarbons
a compound of hydrogen and carbon, such as any of those that are the chief components of petroleum and natural gas.
hydrates
a compound, typically a crystalline one, in which water molecules are chemically bound to another compound or an element.
hydronium ion
a hydrated hydrogen ion, H3O+
acid
a molecule or other entity that can donate a proton or accept an electron pair in reactions.
anion
a negatively charged ion, i.e., one that would be attracted to the anode in electrolysis
oxidation number
a number assigned to an element in chemical combination that represents the number of electrons lost (or gained, if the number is negative) by an atom of that element in the compound
cation
a positively charged ion, i.e., one that would be attracted to the cathode in electrolysis.
Octet rule exceptions
a. molecules that contain an odd number of valence electrons. • This results in less than an octet (8) valence electrons and greater than an octet bonding electrons. • Example : two nitrogen oxides NO an NO2 contain an odd number of valence electrons. • It is impossible to write Lewis dot structure b. Molecules with less than an Octet - around the central atoms are said to be electron deficient and are represented by the hydrides of beryllium. c. Molecules with greater than an Octet - at the third period and beyond of the periodic table d orbitals can become involved in bonding. •This creates the possibility of more than eight valence electrons. •Involvement of on d orbital leads to 10 valence electrons while two d orbitals corresponds to 12 valence electrons. Ex. PF5, PF6
total ionization
all of the ionizable hydrogen is removed from the acid molecule
monoprotic acid
an acid that can donate only one proton, while polyprotic acid can donate more than one proton. Similarly, a monoprotic base can only accept one proton, while a polyprotic base can accept more than one proton.
oxyanion
an anion containing one or more oxygen atoms bonded to another element (as in the sulfate and carbonate ions).
ion
an atom or molecule with a net electric charge due to the loss or gain of one or more electrons.
valence electrons
an electron that is associated with an atom, and that can participate in the formation of a chemical bond; in a single covalent bond, both atoms in the bond contribute one valence electron in order to form a shared pair.
oxyacid
an inorganic acid whose molecules contain oxygen, such as sulfuric or nitric acid
monatomic ion
an ion consisting of a single atom. If an ion contains more than one atom, even if these atoms are of the same element, it is called a polyatomic ion. For example, calcium carbonate consists of the monatomic ion Ca2+ and the polyatomic ion CO32−.
carboxylic acid
an organic acid containing a carboxyl group. The simplest examples are methanoic (or formic) acid and ethanoic (or acetic) acid.
ether
an organic compound in which two alkyl groups are bonded to the same oxygen, having the general formula R-O-R'.
Sketch the wedge
and-dash diagram. This should match the mental picture you formed in step 3.
electron pair, angles and geometry
angle formed by any two electron pairs in a molecule or ion and the central atom between the them and the arrangement of electron pairs around a central atom in a molecule or ion.
electron pair angle
angle formed by any two electron pairs in a molecule or ion and the central atom between the them.
alkane
any of the series of saturated hydrocarbons including methane, ethane, propane, and higher members.
bent
any unshared pairs of electrons will cause this shape usually in group 6 and 16 elements bond angle 104.5 and 109.5 degrees ex. H2O
polar bonds
asymmetrical distribution of electric charge
elements that form diatomic molecule
at common temperatures seven elements form diatomic molecules: Hydrogen (H), Nitrogen (N2), Oxygen (O2), Fluorine (F2), Chlorine (Cl2), Iodine (I2), Bromine (Br2).
Octet Rule
atoms tend to gain, lose or share electrons so as to have 8 electrons •C would like to gain, share electrons •N would like to gain 3 electrons •O would like to gain 2 electrons
effect of multiple bonds on molecular geometry
behave as a single electron pair in establishing molecular geometry.
covalent bond
chemical bond between two atoms that share a pair of electrons.
ionic bonds
chemical bonds arising from attraction forces between oppositely charged ions in an ionic compound
writing chemical equations
chemical equations are always written 'left side' turns into 'right side'. ex Write an equation that shows the formation of a rubidium ion from a neutral rubidium atom Since rubidium is a metal in group 1A it is expected to form a positive ion (a cation) with the magnitude of the charge equal to the group number. The charge is therefore +1. The symbol for the rubidium ion is therefore Rb+ . Rb has an atomic number of 37, with 37 protons and 37 electrons in the neutral atom. In order for the ion to carry a charge of +1 it must have 1 more proton than it does electrons, so the atom must lose 1 electron. The equation that shows the ion forming from the neutral atom is therefore: Rb to Rb+ + e- The neutral Rb atom turns into a Rb+ ion and 1 electron.
double and triple bonds
common in C, N, P, O, S, structures
triple bond
composed of three electron pairs shared between two atoms and are represented as six dots or as three lines.
ionic compounds
compounds where ions are are held by ionic bonds.
overlap of atomic orbitals
concentration of orbitals on adjacent atoms in the same regions of space. Orbital overlap can lead to bond formation.
binary molecular compounds
contains two elements, both nonmetals or one is a metalloid.
ionize
convert (an atom, molecule, or substance) into an ion or ions, typically by removing one or more electrons.
single bond
covalent chemical bond formed by the sharing of one pair of electrons between two bonded atoms
triple bond
covalent chemical bond formed by the sharing of three pairs of electrons between two bonded atoms
double bond
covalent chemical bond formed by the sharing of two pairs of electrons between two bonded atoms
electron pair geometry
description of the distribution of bonding and unshared electron pairs around a bonded item.
electron affinity
discrete energy needed to add an electron
procedure for predicting molecular geometries
draw the Lewis diagram, count electron pairs around the central atom, both bonding and lone pairs, determine electron pair and molecular geometries. This is best done by reason rather than by memorization Ask yourself, and picture the answer in your mind: "where will the electron pairs go to be as far apart as possible?"
isomers
each of two or more compounds with the same formula but a different arrangement of atoms in the molecule and different properties.
Two electron pairs
electron pair and molecular geometries are both linear, bond angle is 180 degrees.
Three electron pairs
electron-pair geometry is trigonal planar. Bond angles are 120 degrees (i) all electron pairs bonding: molecular geometry is trigonal planar. (ii) Two electron pairs bonding, one lone pair: Molecular geometry is angular (bent).
octet rule
elements of low atomic number tend to lose, gain or share electrons in such a way that they each have eight electrons in their valence shells, giving them the same electronic configuration as a noble gas.
resonance
elements with multiple lewis structures
enthapy
energy necessary to break a bond
neutral to ion
ex: potassium, cobalt, bromine, oxygen Since potassium is a metal in group 1A it is expected to form a positive ion (a cation) with the magnitude of the charge equal to the group number. The charge is therefore +1. The symbol for the potassium ion is therefore K+ . K has an atomic number of 19, with 19 protons and 19 electrons in the neutral atom. In order for the ion to carry a charge of +1 it must have 1 more proton than it does electrons, so the atom must lose 1 electron. The equation that shows the ion forming from the neutral atom is therefore: K K+ + e- Note that chemical equations are always written 'left side' turns into 'right side'. The neutral K atom turns into a K+ ion and 1 electron. cobalt: Since cobalt is a transition metal, it is not possible to predict its ion charge from its group number. The charge on the ion is given by the roman numeral in its name. The charge is therefore +3. The symbol for the cobalt(III) ion is therefore Co3+ . Co has an atomic number of 27, with 27 protons and 27 electrons in the neutral atom. In order for the ion to carry a charge of +3 it must have 3 more protons than it does electrons, so the atom must lose 3 electrons. The equation that shows the ion forming from the neutral atom is therefore: Co Co3+ + 3 e- Note that chemical equations are always written 'left side' turns into 'right side'. The neutral Co atom turns into a Co3+ ion and 3 electrons. bromine: The element bromine, Br, is a Group 7A nonmetal. Nonmetals form anions in ionic compounds. Anion Name : Remove ending, add 'ide' = bromide ion Anion Charge : group # - 8 = 7 - 8 = -1 Anion Symbol : Br- Br has an atomic number of 35, with 35 protons and 35 electrons in the neutral atom. In order for the ion to carry a charge of -1 it must have 1 more electron than it does protons, so the atom must gain 1 electron. The equation that shows the ion forming from the neutral atom is therefore: Br + e- Br- Oxygen: The element oxygen, O, is a Group 6A nonmetal. Nonmetals form anions in ionic compounds. Anion Name : Remove ending, add 'ide' = oxide ion Anion Charge : group # - 8 = 6 - 8 = -2 Anion Symbol : O2- O has an atomic number of 8, with 8 protons and 8 electrons in the neutral atom. In order for the ion to carry a charge of -2 it must have 2 more electrons than it does protons, so the atom must gain 2 electrons. The equation that shows the ion forming from the neutral atom is therefore: O + 2 e- O2-
molecules with fewer that four electron pairs around the central atom
fluorides of beryllium and boron. These atoms are surrounded by two and three pairs of electrons, respectively.
metallic bond
forces of attraction between delocalized electrons in a metallic crystal and the metal ions.
ions of transition metals
forms one or more stable ions which have incompletely filled d orbitals.
octet rule
general rule that atoms tend to form stable bonds by sharing or transferring electrons until the atom is surrounded by a total of eight electrons
chemical bond
general term that sometimes includes all of the electrostatic attractions among atoms, molecules, and ions, but more often refers to covalent and ionic bonds.
crystals
geometric ion arrangement in ionic bonds
main group elements
group of elements whose lightest members are represented by helium, lithium,beryllium, boron, carbon, nitrogen, oxygen, and fluorine as arranged in the periodic table of the elements.
chromate ion -
has a molecular formula of CrO42
permanganate ion
has a molecular formula of MnO4
polar molecule
has a net dipole as a result of the opposing charges (i.e. having partial positive and partial negative charges) from polar bonds arranged asymmetrically. Water (H2O) is an example of a polar molecule since it has a slight positive charge on one side and a slight negative charge on the other.
hydrogen
has one electron and wants to look like helium (octet)
triprotic
has three ionizable hydrogens. such as phosphoric acid (H3PO4) and citric acid (C6H8O7),
isoelectronic
have the same numbers of electrons or have the same electronic structure.
Isoelectronic
having the same (iso) number of electrons.
isoelectronic
having the same electron configuration
ionization energy
how much energy it takes to pull an electron away Coulomb's law
naming hydrates
hydrate, anhydrous compound, name of anhydrous compound, # of water molecules, name of ionic compound ex: CuSO4 . 5H2O, CuSO4, Copper (II) Sulfate, 5 (penta) = copper(II)sulfate pentahydrate
hydrated proton
hydronium ion
prefixes used in the naming of molecular compounds
if an elemental name has no prefix in a binary molecular compound, assume that there's only one atom of that element in the molecule. Prefixes: 1=mono, 2= di, 3= tri, 4=tetra, 5=penta, 6=hexa, 7=hepta, 8=octa, 9=nona, 10=deca. The letter o at the end of the prefix mono and the letter a in the prefixes for four to ten are omitted if the resulting word sounds better.
prefixes and suffixes in the anions of chlorine
if the number of oxygens is one larger than the number in the ic acid, the prefix per- is placed before both the acid and the anion names. if the number of oxygens is one smaller than the number in the -ic acid, the suffixes -ic and -ate are replaced with -ous and -ite. if the number of oxygens is two smaller than the number in the -ic acid (one smaller than the number in the -ous acid) the prefix hypo- is placed before both the acid and anion names and the -ous and -ite suffixes are kept. the name of an acid with no oxygen is hydro- followed by the name of the nonmetal changed to end in -ic. The monoatomic anion from the acid is named by the rule for monoatomic anions by which the elemental name is changed to end in -ide.
cyanide ion
in inorganic cyanides, such as sodium cyanide, NaCN, this group is present as the negatively charged polyatomic cyanide ion (CN−)
core electrons
inside electrons (nearest the nucleus) shields outer electrons
dichromate ion
ion with a 2 charge used as an oxidizing agent
chemical bonds (3 types)
ionic covalent metallic
monoatomic ions
ions consisting of one atom of a single element (usually in the same family) K+ Na+ Cl- N3- Ca2+
helium
is a noble gas that does not have an octet (8 electrons in the outermost shell)
acetate ion
is acetic C2H3O2
bond angle
is the angle formed between three atoms across at least two bonds.
central atom
is the least electronegative element that isn't hydrogen. Connect the outer atoms to it by single bonds.
3D molecular structures
linear planar tetrahedral
Electron Pair Geometry for 2 electron groups
linear - e-pair domain geometry linear - e-molecular geometry 2 electron pairs around center/atom 2 bonded pairs 0 lone pairs 180 degree bond angle ex: BeH2
geometry arising from two electron pairs, two bonded atoms
linear, 180 degree bond angle.
metallic bonds
mainly occurs between metals holds atoms of a metal together • Most metals have one, two or three valence electrons that are loosely held. They can be easily removed to form ions. • When they are surrounded by other metals, the valence electrons of the metals freely move around according to the electron-sea model while the cations remain fixed. • Because electrons are free to move around this allows metals to have a high electrical conductivity. • A metallic bond occurs due to the attractive forces between positively charged metal ions and the negatively charged valence electrons of neighboring metallic ions. • These electrons are delocalized, spread out, unlike the electrons in the covalent bond that are localized between the specific atoms. • Electron cloud around atoms • Good conductors at all states, lustrous, very high melting points • Examples; Na, Fe, Al, Au, Co
covalent bonds
mainly occurs between nonmetals some valence electrons shared between atoms
acid
molecular compound that reacts readily with water to produce hydrated hydrogen ion (proton) and the anion Examples of acids: • HCl • HNO3
isoelectronic H and Li
monoatomic hydride and lithium ions, H- and Li+, dubplicate electron configuration of helium with just two electrons: 1s2
Fluorine
most electronegative element
naming of monoatomic cations
named by writing the name of the element ex: Group 3A metal Aluminum, Sodium, Calcium Al to Al3+ +3e-
anion
negatively charged ion
isoelectronic with neon
nitrogen oxygen fluorine sodium magnesium aluminum
covalent bond types
nonpolar polar
valence electrons
outermost electrons
ionization energy
outermost energy requires the least energy to remove it, energy needs increase as the orbital becomes closer to the nucleus.
lone pairs
pair of valence electrons in a molecule that are not used for bonding.
non-polar bonds
pertaining to a bond or molecule having a symmetrical distribution of electric charge.
isoelectronic with argon
phosphorus sulfur chlorine potassium calcium scandium
polyatomic
phosphorus,(P4) sulfur (S8) and carbon(buckyball) C60 . • However, in most instances these elements are written without subscripts.
crystal
poor conductors of electricity, high temps are required to break ionic bonds, free ions from one another and melt the crystal to a liquid. Common white crystals are salt and sugar. Based on the microscopic arrangement of atoms inside it, called the crystal structure. A crystal is a solid where the atoms form a periodic arrangement.
cation
positively charged ion
lewis structure
representations of molecules showing all electrons, bonding and nonbonding.
electronegativity
scale of the relative ability of an atom of one element to attract the electron pair that forms a single covalent bond with an atom of another element
isoelectronic with krypton
selenium bromine rubidium strontium yttrium
d orbitals and f orbitals
shielded under s and p orbitals d orbitals hide under s orbitals above them
Lewis symbol
simplified Bohr diagrams which only display electrons in the outermost energy level. The omitted electrons are those in filled energy levels, which do not contribute to the chemical properties of the species in question. Useful for studying elemental properties and reactions. Constructed by placing dots representing electrons in the outer energy around the symbol for the element. For many common elements, the number of dots corresponds to the element's group number.
molecule
smallest unit particle of a pure substance that can exist independently and possess the identity of the substance.
alloy
solid mixture of two more more elements that has macroscopic metallic properties.
multiple bonds energy and length
stronger and shorter than single bonds. Bond enthalpies measure the energy necessary to break a bond. single = 154 (bond length), 346 = enthalpy double = 134 (bond length), 612 = enthalpy triple = 124 (bond length), 835 = enthalpy
diprotic
sulfuric acid (H2SO4), carbonic acid (H2CO3), hydrogen sulfide (H2S), chromic acid (H2CrO4), and oxalic acid (H2C2O4) have two acidic hydrogen atoms
beryllium and boron
tend to form covalent bonds by sharing electrons, not forming ions
geometry arising from four electron pairs, four bonded atoms
tetrahedral
geometry arising from four electron pairs, two bonded atoms
tetrahedral, bent/angular
geometry arising from four electron pairs , three bonded atoms
tetrahedral, pyramidal, triangular, "squashed down" pyramid
geometry arising from four electron pairs, three bonded atoms
tetrahedral, triangular, "squashed down" pyramid, pyramidal
transition elements that form only one ion
the charge is included in the name of an ion only when the ions of an element exhibit more than one common charge. Zinc and silver form ions of only one charge under ordinary chemical conditions. Charge is not included in the names of these ions. Memorize these charges as they are not indicated by the periodic table.
organic chemistry
the chemistry of carbon compounds (other than simple salts such as carbonates, oxides, and carbides).
method of naming compounds composed of two nonmetals
the first word is the name of the element appearing first in the chemical formula, including a prefix to indicate the number of atoms of that element in the molecule. Second word is the name of the element appearing second in the chemical formula, changed to end in ide, and also including a prefix to indicate the number of atoms in the molecule.
difference between the formula of an element and an ion of that element
the formula is the smallest representation of a substance using symbols and numbers, while the ion is the charged particle of some element or molecule.
naming of oxyanions derived from oxyacids with names ending in "ic"
the formula of the anion is the formula of the acid without the hydrogens(s). there is a negative charge equal to the number of ionizable hydrogens in the acid, the name of the anion is the name of the central element of the acid changed to end in ate.
hydronium ion
the ion H3O+, consisting of a protonated water molecule and present in all aqueous acids
Lewis symbol
the symbol for an element showing a dot for each valence electron in the element or ion.
formula unit
the tiniest individual unit in a sample of the element is one atom
molecular geometry
three-dimensional arrangement of the atoms that constitute a molecule. It determines several properties of a substance including its reactivity, polarity, phase of matter, color, magnetism, and biological activity.
resonance structures
two forms of a molecule where the chemical connectivity is the same but the electrons are distributed differently around the structure. Resonance occurs when electrons can flow through neighboring pi systems.
anions formed by the progressive ionization of polyprotic acids
two or more ionizable hydrogens
electrical conductivity of solid ionic compounds
typically have high melting and boiling points, and are hard and brittle. As solids they are almost always electrically insulating, but when melted or dissolved they become highly conductive, because the ions are mobilized.
trigonal pyramidal
usually seen in group 5 and 15 elements bond angle is 107 degrees there's an unshared pair of electrons ex. NF3, PF3
electrical conductivity of liquid ionic compound
when an ionic compound is heated until the solid changes to the liquid state, it conducts electric current.
process for writing the formula of an ionic compound
write the formula of the cation, followed by the formula of the anion, omitting the charges. Insert subscripts to show the number of each ion needed in the formula unit to make the sum of the charges equal to zero with the fewest number of ions possible. if only one ion is needed, omit the subscript. If a polyatomic ion is needed more than once, enclose the formula of the ion in the parentheses and place the subscript after closing the parenthesis.
naming of ionic compounds
write the name of the cation, write the name of the anion
transition metals that only form one ion
zinc (Zn) and silver (Ag) Zn2+ Ag+
electron pair repulsion
• Electron pairs in the bond repel other electron pairs giving rise to a repulsion force that affects the overall energy of the molecule. Why do the electron pairs repel each other? • Because the lower or more negative the energy, the more stable the molecule. • In order to minimize the repulsions and to maximize the negative energy content, the electron pairs position themselves so as to minimize electron pair repulsion. • The bonds (made up of electrons) and electron pairs stay as far away from each other as possible
compounds made from 2 nonmetals
• Elements above and to the right of the stair step (Darker line) in the periodic table are nonmetals. • Because of the electronegativity differences between non metals, the bonds formed between them are not ionic but are polar covalent. • As such the smallest unit of such a compound made up of non metals is called is a molecule. These compounds are referred to as binary molecular compounds which means they are composed of two different elements joined together by polar covalent bonds.
bond polarity
• HCl is POLAR because it has a positive end and a negative end due to difference in electronegativity Cl has a greater share in bonding electrons than does H because it is more electronegative. Cl has slight negative charge (-d) because it is electronegative and H has slight positive charge (+d)
cation
• If the atoms loses electrons it forms a positively charged ion called a cation.
Ions that are isoelectronic with Neon
• Ne configuration • Non metals • N3- configuration • O2- configuration • F- configuration • Metals • Na+ configuration • Mg2+ configuration • Al3+ configuration 1s22s22p6
Monoatomic vs. Polyatomic ions
• Polyatomic ions which consists of more than one (poly) element in one ion example • Phosphate ion PO43- • Carbonate ion CO32- • Sulfate ion SO42- • Nitrate ion NO3- • These polyatomic ions being negatively charged form ionic bonds with metals • Ca3 (PO4)2 • Na3 PO4 • Na2 SO4
molecular geometry
• Properties of molecular substances depend on the structure of the molecule. • The structure is influenced by many factors, such as - the skeletal arrangement of the atoms - the kind of bonding between the atoms • ionic, polar covalent, or covalent - the shape of the molecule • Bonding theory should allow you to predict the shapes of molecules.
electronegativity trends
• The ability of an atom of that element in a covalent bond to attract the shared/bonded electrons to itself. • Fluorine is the most electronegative element. • Trends across the periodic table. • Electronegativity increases from left to right • Across the periodic table • Electronegativity decreases from top to bottom, down the group of the periodic table.
electron pair geometry
• The electron pair geometry takes into account ALL the electron pair groups around the central atom regardless of whether the electron pairs are bonded pair electrons or unbonded pairs (lone pairs)
shape
• The molecular geometry of a molecule may be determined from the molecular formula by drawing the Lewis Dot Structure, counting the number of electron pair groups around the central atom and invoking the electron pair repulsion mode.
crystalline ion structure
• The ratio of the ions that are needed to form the crystalline structure depends on the charges on the ions that are combining together in the ionic bond. • The goal is to make neutral crystals • Example: • NaCl - requires 1 sodium ion to 1 chloride ion • CaCl2- requires 1 calcium ion to 2 chloride ions • AlCl3- requires 1 calcium ion to 3 chloride ions
periodic table trends
• Two rules which summarize these trends are : • Metals lose their valence electrons to form cations with a charge that is the same as the group number. • Nonmetals gain enough valence electrons to complete their octet and form anions with a charge that is the difference between their group number and the number eight.
anion
•If the atom gains electrons it forms a negatively charged ion called an anion.
formulas of ionic compounds
•Ionic compounds are a combination of a cation, with a positive charge, and an anion, with a negative charge. In order to have no net charge, the positive and negative charges must balance. If the cation and anion have the same magnitude of charge, the condition for zero net charge is met. However, if the magnitudes of the charges are different, the subscripts on the cation and anions are taken so that the net charge is zero. 1. Write down the formulas for the cation and anion, keeping in mind their charges, without explicitly writing them down. 2. Insert as subscripts, whole small numbers, that correspond to the magnitude of the charge on the opposite ion. The subscripts for the cation comes from the charge on the anion. The subscript for the anion comes from the charge on the cations. Subscripts of one are not explicitly written. If an ion is polyatomic, it is endorsed in a set of parentheses and the subscript is written outside the closing parenthesis. ex: sodium oxide = Na+O2-, Na2O
ions formed by one element
•Transition elements - can form cations by losing "s" electrons as well as "d" electrons to achieve multiple oxidation states. •These cations are named by writing the name of the element followed by the ionic charge, written as a Roman numeral in parenthesis. The number of electrons lost or gained by an atom of a representative element depends on its position relative to the nearest inert gas. Metals will lose enough electrons to achieve an octet in the outer shell while nonmetals will gain enough electrons to achieve an octet in the outer shell. ex: Iron Fe2+, Fe3+ cation Iron (II) = cation name Iron (III)
Formation of Ions within the group (anions)
● Negative ions generally form in group VA, VIA & VIIA elements when electrons are gained. ● The gaining of electrons onto a neutral atom results in an excess of negative ions ● The general trend is as follows: ● There is an excess of negative ions • Group 7 non metals (gains 1 electron) • Group 6 non metals (gains 2 electron) • Group 5 non metals (gains 3 electron) ion 1- ion 2- ion 3-
Formation of Ions within the group (cations)
● Positive ions generally form in group IA,IIA & IIIA elements when electrons are lost. ● The loss of electrons from a neutral atom leaves behind an excess of positive ions ● The general trend is as follows: • Group I A metals ( loses 1 electron) • Group 2A metals (loses 2 electrons) • Group 3A metals (loses 3 electrons) ion 1+ ion 2+ ion 3+