Module 1: Atomic Structure & Chemical Properties PDF

Module 1. How does atomic structure have an effect on the properties of chemical elements? Matter and sustainability 1.1 Review: Subatomic particles, Z, A. History of the atom: theories and models ATOM means “indivisible” Atom Atoms are the smallest parts in which an element can be divided, without losing its properties and identity. Atoms consist in: - Protons - Neutrons - Electrons The number of these particles determine the type of element and other properties of the atom. Fact: The modern picture of the atom includes many more subatomic particles, but these three particles are enough to explain chemical and physical properties of elements. Subatomic particles Electrons Protons Neutrons Are located outside the nucleus Are located inside the of an atom. Are located inside the nucleus of an atom. They carry a -1 electrical nucleus of an atom. They carry no electrical charge and have an almost They carry a +1 electrical charge (0) and have a mass of 1/1836 atomic mass charge and have a mass of 1 mass of 1 atomic mass units. atomic mass unit (amu). unit (amu) They move rapidly around the heavy nucleus Protons and electrons carry equal but opposite electrical charges, so a neutral atom (one with no electrical charge), must have the same number of protons as of electrons, in order to balance each other’s charges. What makes atoms different? It’s all about QUANTITY of subatomic particles they have. All the atoms of the same element always have the same number of protons. This distinguishes each element (as an id). The number of protons never changes for an element. Atomic mass & Atomic number The atomic number of an atom (Z) is equal to the number of protons in the nucleus of the atom. The mass number of an atom (A) is equal to the sum of protons & neutrons in the nucleus of the atom. Both are whole numbers. The chemists use the Nuclear Symbol to describe the nucleus of the atom. X: chemical symbol of the element A: mass number (#p + #n) Z: atomic number (#p) which equals the number of electrons in a neutral atom Here is an example of a nuclear symbol: Periodic table: Atomic number is written as an integer/whole above the symbol. Mass number is not written. Instead, atomic weight is written as a decimal below the symbol. Learning check Learning check Isotopes Are atoms of an element with the same number of protons in the nucleus but different number of neutrons (have the same atomic number but different mass number). Most of the elements exist as a mixture of their isotopes… Isotopes in nature Example: chlorine is a mixture of two isotopes, chlorine-35 and chlorine-37. Example: naturally occurring magnesium consists of three isotopes with same protons, but different number of neutrons. The periodic table’s atomic weight does not represent the mass number for an individual atom. It shows an average of the mass number for a random sample of atoms (isotopes). Therefore, is a fractional number. Ion It is an atom that has lost or gained electrons, so it is no neutral anymore, it has an electric charge (+ or -) Cation when an atom loses an electron, it is positively charged. protons(+) → electrons(-) = (+) charge Anion when an atom gains an electron, it is negatively charged. protons(+) ← electrons(-) = (-) charge Videos https://www.youtube.com/watch?v=_lNF3_30lUE https://www.youtube.com/watch?v=GkfZnY1Rn2I https://www.youtube.com/watch?v=LhveTGblGHY 1.2 Electronic configuration Colors of aurora borealis Why is color emitted after the collision between solar wind and atmospheric particles? This happens because atoms, when excited and then return to their ground state, emit photons, which are the particles that make up light. Electrons in ions become excited upon impact from solar wind particles, and they jump to a higher energy level. This makes the electron unstable, and this state lasts for a short time. When the electron returns to its ground state, it releases energy in the form of light. Quantum Mechanical Model As electrons are always moving around the nucleus, their location cannot be determined exactly, only as a probability zone or space which is called orbital. Location and energy of electrons is speciﬁed as an address, using three terms: level(shell), sublevel (subshell), and orbital. Energy levels/shells Each atom has 7 energy levels or shells, represented in orbits around the nucleus in which electrons orbit. if you pay attention to your periodic table, you can also identify 7 periods corresponding to these seven energy levels. To identify each level, a quantum number called n was assigned, ranging from n=1 to n=7 Designated from lower to higher energy level. The electrons are arranged in different energy levels depending on the period in which they are located in the periodic table. Example: Sodium (Na), which is in period 3, its electrons are distributed over 3 energy levels. Bromine (Br), which is in period 4, its electrons will be distributed over 4 levels. Francium (Fr), which is in the last period of the periodic table, will have its electrons distributed over 7 energy levels. Energy sublevels/subshells Each of the 7 levels also has energy sublevels, which can have up to 4 sublevels. These sublevels are designated with the letters S, P, D, F. Each sublevel has the same shape but different spatial orientations. Sublevels have different magnetic orientations, and each orientation can hold a maximum of 2 electrons. Each sublevel (s, p, d and f) has a speciﬁc number of orbitals, as shown in the table to the right. Orbitals’ capacity Letter of Number of Number of e- subshell orbitals max. Examples s 1 2 p 3 6 d 5 10 f 7 14 Electronic configuration An electronic conﬁguration is a representation in detail of the arrangement of electrons in atoms. The way electrons can be arranged is deﬁned by three rules or principles: - Aufbau’s principle - Pauli’s exclusion principle - Hund’s Rule of Maximum Multiplicity Aufbau’s principle Also known as the electronic distribution principle. It states that each electron occupies the lowest available energy orbital. Electrons are added to atomic orbitals in order of low energy. Therefore, the ﬁrst step to determine the electronic conﬁguration of an element is to know the sequence of atomic orbitals. Pauli’s exclusion principle The principle states that a maximum of two electrons can occupy the same atomic orbital, but only if they have opposite spins. Electron spin is represented with arrows: An upward arrow (↑) represents spin in a positive direction, while a downward arrow (↓) represents spin in the opposite direction. Hund’s Rule of Maximum Multiplicity The fact that negatively charged electrons repel each other affects their distribution. The rule states that electrons with the same spin must occupy different orbitals before electrons with opposite spins can occupy the same orbitals. Write electronic conﬁgurations for the following ﬁve elements: - Fluorine (F) e= - Sulfur (S) e= - Iron (Fe) e= - Gallium (Ga) e= - Strontium (Sr) e= Valence Shell & Electrons Valence shell of an atom is the outermost occupied shell that contains electrons (total or partially), including all its occupied subshells. Valence electrons are the total number of electrons in the valence. shell. 2 2 4 Ex.: 1s 2s 2p (valence shell is 2 and v. electrons 6). Valence Shell & Electrons Periodic Table: Elements with the same valence shell belong in the same period of the periodic table. Elements with same valence electrons are members of the same group and have similar chemical properties.. Valence Shell Example The electronic conﬁguration of an atom of Tin - Sn, atomic number 50 is as follows: The Valence shell of an atom is the highest coefﬁcient used in the electronic conﬁguration Valence shell of Tin: ______________ The Valence electrons is the addition of the electrons in all the subshells of the valence shell Valence electrons of Tin: _____________ Valence Shell Example The electronic conﬁguration of an atom of Tin - Sn, atomic number 50 is as follows: The Valence shell of an atom is the highest coefﬁcient used in the electronic conﬁguration, so for Sn is 5s2 5p2 (n=5). Therefore, Sn is located in period 5 of the Periodic Table. The Valence electrons is the addition of the electrons in all the subshells of the valence shell, so for Sn is v.e- = 2+2 = 4. Therefore, Sn is in group 4A of the Periodic Table. This part is only applicable for Representative Elements (groups 1A to 8A). Blocks of elements by orbital The last orbital occupied in an atom (last letter of el. conﬁguration) deﬁnes four blocks of elements in the periodic table: s, p, d, and f. Representatives →last orbital occupied is s or p. Transition → last orbital occupied is d. Inner-trans. → last orbital occupied is f.. Last orbital occupied NOT always matches the valence shell Electronic conﬁgurations & the periodic table The electronic conﬁgurations can be used to locate an element in the periodic table and obtain information about it. Also, the periodic table can be used to write the electronic conﬁguration of a speciﬁc element by reading the blocks from left to right across the periods. Examples: Write the electronic conﬁguration of Al, underline valence shell 13 Electronic conﬁgurations & the periodic table The electronic conﬁgurations can be used to locate an element in the periodic table and obtain information about it. Also, the periodic table can be used to write the electronic conﬁguration of a speciﬁc element by reading the blocks from left to right across the periods. Examples: Write the electronic conﬁguration of Al, underline valence shel 13 2 2 6 2 1 Al, atomic number 13, it’s in 3p block, then: 1s 2s 2p 3s 3p (period 3, group 3A → shell 3, 3 valence electrons.) Electronic conﬁgurations & the periodic table Obtain the information about P: 15 Electron conﬁguration: Highest energy level (valence shell): Valence Electrons: Period: Group: Electronic conﬁgurations & the periodic table Obtain the information about P: 15 2 2 6 2 3 Electron conﬁguration: 1s 2s 2p 3s 3p Highest energy level (valence shell*): 3 Valence Electrons: 5 Period: 3 Group: 5A** * Highest energy level is the same as valence shell ** Group is valid only for Representative Elements Electronic conﬁgurations & the periodic table For the electronic conﬁgurations of the elements you did before identify their: - Valence shell (highest coefﬁcient used, including all its subshells) - Valence electrons (addition of all electrons in the valence shell) - Block of P. Table (last orbital s, p, d or f occupied by an electron) Fluorine, F, 9 electrons Sulfur, S, 16 electrons Iron, Fe, 26 electrons Gallium, Ga, 31 electrons Strontium, Sr, 38 electrons Electronic conﬁgurations & the periodic table For the electronic conﬁgurations of the elements you did before identify their: Valence shell Valence Last orbital occupied Period Group coeff. electrons Fluorine, F, 9 electrons Sulfur, S, 16 electrons Iron, Fe, 26 electrons Gallium, Ga, 31 electrons Strontium, Sr, 38 electrons Noble Gas Conﬁgurations Noble gas conﬁgurations: all noble gases (group 8A) have electronic conﬁgurations that end with completely ﬁlled s and p subshells of the highest occupied shell. 2 6 Their valence shell is full with 8 electrons (ns np ), except for the smallest one, Helium ( He), whose valence shell ﬁlls with 2 only 2 electrons. Noble Gas Conﬁgurations Noble Gas Electronic conﬁguration 2 He 10 Ne 18 Ar 36 Kr 54 Xe 86 Rn Noble Gas Conﬁgurations Abbreviated Conﬁgurations or Kernel electronic conﬁguration Noble gas conﬁgurations can be used to write electronic conﬁgurations of larger atoms in an abbreviated form, in which the noble gas symbol enclosed in brackets is used to represent all the electrons found in that noble gas conﬁguration. Each noble gas abbreviates conﬁgurations of elements in the period below its own: He (period 1) abbreviates elements in period 2. Abbreviated Conﬁgurations The conﬁguration of a given element is abbreviated with the noble gas in the period directly above its own. Example: Ar (period 3) abbreviates elements in period 4. Abbreviated Conﬁgurations or Kernel electronic conﬁguration Write abbreviated electronic conﬁgurations for the elemen [Symbol of noble gas] Rest of the conﬁguration Fluorine, F, 9 electrons Sulfur, S, 16 electrons Iron, Fe, 26 electrons Gallium, Ga, 31 electrons Strontium, Sr, 38 electrons Energy/orbital diagrams Orbital diagrams are a graphical way to represent orbitals and their electrons. They’re like electronic conﬁgurations ampliﬁed. Each Orbital is represented by a small horizontal line with the orbital name under it and electrons are represented by arrows on it:. Orbital diagrams The orbitals within each subshell ﬁll with electrons in consecutive order, but must obey two rules: 1. Pauli exclusion principle (electrons spin in opposite directions) 2. Hund's rule (electrons occupy all empty orbitals of the subsh. ﬁrst) Orbitals in subshells It’s important to remember the maximum number of orbitals and of electrons per subshell, as follows:. Orbitals in subshells It’s important to remember the maximum number of orbitals and of electrons per subshell, as follows:. Orbital diagram Fluorine, F, 9 electrons Sulfur, S, 16 electrons Iron, Fe, 26 electrons Gallium, Ga, 31 electrons Strontium, Sr, 38 electrons 1.3 Electronic configuration of carbon. Carbon element Most of the materials you see around you are derived from organic compounds, based on carbon. More than 10 million organic compounds are known compared to just 1 million inorganic compounds. Friedrich Wöhler believed that organic compounds were based on the vital force, coming only from living beings, and thus all organic compounds should come from living beings. Wöhler's discovery opened the door to organic chemistry, the chemistry of carbon, and for this reason, he is known as the father of organic chemistry. Electronic conﬁguration and carbon hybridization Carbon is in group IV and period 2 of the periodic table, with an atomic number of 6 and a mass of 12.0107 Having 6 electrons, its ground state electronic conﬁguration would be: Sublevels have different magnetic orientations, and each orientation can hold a maximum of 2 electrons. Electronic conﬁguration and carbon hybridization The organic compounds in which carbon participates are tetravalent, meaning they form 4 bonds, for example in methane (CH4). Promotion of electrons to empty orbitals Orbital hybridization Carbon hybridization: 1.4 Periodic properties Periodic table The periodic table can provide you with valuable information that will help you recognize which compounds can form and which cannot, just as scientists do to create new materials that are useful in your daily life. Tool to learn, understand, and predict the properties of elements. Characteristics of periodic table - Consists of squares or boxes containing the 118 elements recorded to date. - Elements arranged in order of increasing atomic number. Characteristics of periodic table - Organized into 7 periods (horizontal rows) - 18 groups or families (vertical columns) Periods A period is a horizontal row of elements arranged according to increasing atomic numbers. Periods are numbered from top to bottom of the periodic table with natural numbers always. Groups or families A group is a vertical column of elements that have similar chemical properties. Traditional designation uses a natural number or Roman numeral and a letter (A or B) at the top of the column. - Group A: representative elements because they have a wide range of physical and chemical properties - Group B: comprises transition elements, which are located in the center of the periodic table. Modern (but not universally-used) designation uses only a number from 1 to 18. ELEMENTS AND THE PERIODIC TABLE Each element belongs to a group and a period of the periodic table. EXAMPLES OF GROUP AND PERIOD LOCATION FOR ELEMENTS Calcium, Ca, element 20: group IIA (2), period 4 Silver, Ag, element 47: group IB (11), period 5 Sulfur, S, element 16: group VIA (16), period 3 Modern periodic table Elements 58-71 and 90-103 are not placed in their correct periods , but are represented below the main table , to keep the table compact Classiﬁcation of elements. The periodic table can be used to classify elements in several ways: - by block or section of elements as Representative, Transition, or Inner-transition Elements. - by metallic character as Metal, Nonmetal, or Metalloid. - by chemical family as Alkali metal, Alkaline earth metal, Halogen or Noble Gas. Periodic properties. Periodic Properties of elements change in a systematic way within the periodic table, they show trends (change in a predictable way). Periodic Properties: Valence electrons Metallic character Atomic size (radius) Electronegativity Ionization energy (potential) Reactivity Valence electrons The chemical properties of representative elements are mostly due to the Increasing valence electrons valence electrons, which are the electrons in the Increasing valence shell outermost energy level. The group number is equal to the number of valence electrons, which occupy the s and p orbitals, increasing to the right. All elements in group 1A have one valence electron in an s orbital. Metallic character trend INCREASING METALLIC CHARACTER Elements in the same period of the periodic table become more metallic or less nonmetallic from right to left across the period. LESS NON METALLIC Elements in the same group of the METALS periodic table become more metallic or less nonmetallic from top to bottom down the group. METALS MORE METALLIC Metallic character Atomic size (radius) trend The distance between the nucleus and INCREASING ATOMIC SIZE valence electrons (farthest) determines the atomic radius. When you move to the right on a period, increasing the atomic number, the force of attraction from the - nucleus to e increases (as it has more SMALLER protons), making the atomic radius shorter. ATOMS For representative elements in the same period, atomic size or radius (of atom’s sphere) increases from right to left in the period and from top to bottom down the group. BIGGER ATOMS 1A 2A 3A 4A 5A 6A 7A Group Period Increasing atomic size 1 A bigger nucleus attracts more the electrons, then 2 the atomic size gets smaller. 3 Increasing atomic size (# v. sh = period) Atoms in the same group 4 have same number of valence electrons but the valence shell is bigger as 5 the period increases. 6 Atoms enlarged ~60 million times. Atomic radius in picometers (10-12) INCREASING ELECTRONEGATIVITY Electronegativity trend Electronegativity is the ability of an atom to attract bonding electrons of another atom to itself (disposition to react chemically). Hydrogen is an exception to the NonMetals: For representative elements the general electroneg. trend More trend is an increase from bottom to top up Electronegative the group and left to right across the period. Nonmetals have high electronegativity compared to metals. Noble gases don’t have electronegativity Metals: Less (don’t form bonds easily). Electronegative Ionization energy trend or INCREASING IONIZATION ENERGY ionization potential trend Ionization energy is the force needed to pull away or remove an electron from the outermost energy level. Since electrons that are closer to the nucleus are more attracted to it, more force has to be used to NonMetals: remove them as the atom size is smaller. Higher ionization Ionization energy trend is opposite to atomic size energy trend. For representative elements, the ionization energy or potential trend increases from bottom to top up the group and left to right across the period. Metals: Nonmetals have high ionization energy lower compared to metals. ionization energy Reactivity trend Reactivity is a measure of how easily an atom reacts with another. This property Nonmetals’ works different for metals and non-met. reactivity increases Metals react by losing electrons, so their reactivity increases for elements whose atoms do not attract their outermost electrons strongly (bigger atomic radius, lower ionization energy) Non-metals react by gaining electrons, so their reactivity increases for elements whose atoms attract other electrons Metals’ strongly (smaller atomic radius, higher reactivity electronegativity) increases Summary of periodic properties Octet Rule An octet is an atom that contains 8 valence electrons (8 electrons in the last shell). OCTET RULE: Atoms will gain or lose sufﬁcient electrons to achieve an outer shell electron arrangement identical to that of a noble gas. This arrangement usually consists of eight electrons in the valence shell, except for He. Octet Atoms acquire octets: By forming compounds (chemically bonding with other atoms). To become more stable. By losing, gaining, or sharing valence electrons. Octets are often represented by electron-dot symbols, where the dots represent the valence electrons. To acquire octets, the atoms often transform in ions. 1.5 Types of bonds Ion Ion is an atom that is not neutral anymore. Now it has an electrical charge (+ or -) because of a loss or gain of electrons. Formation of positive ions (cations) Lost e- Metals form: - Octets by losing all of their valence electrons. - Form positive ions with the electron conﬁguration of the nearest noble gas (previous period). - Positive ions (cations) have fewer electrons than protons, then have positive ionic charge. 1+ - 1+ Group 1A metals → lose 1 valence e → ion Group 2A metals → lose 2 valence e- → ion 2+ Group 3A metals → lose 3 valence e- → ion 3+ + Positive ion of sodium (Na ) Sodium 11Na - Is a metal located in group 1A. Conﬁguration 1s22s22p63s1 - Achieves an octet by losing its 1 valence electron. 1+ + - Forms a positive ion with an ionic charge of 1+ (Na or Na ) 2 2 6 - The ion’s new elec. conﬁguration is 1s 2s 2p (as noble gas 10Ne) Sodium Sodium atom: ion: 11p+ 11p+ 11e- 10e- 0 1+ Neutral Positive charge charge + Positive ion of magnesium (Mg ) Magnesium 12 Mg - Is a metal located in group 2A. Conﬁguration 1s22s22p63s2 - Achieves an octet by losing its 2 valence electrons. 2+ - Forms a positive ion with an ionic charge of 2+ (Mg ) - The ion’s new elec. conﬁguration is 1s22s22p6 (as noble gas 10Ne) Magnesium Magnesium atom: ion: 12p+ 12p+ - - 12e 10e 0 2+ Neutral Positive charge charge Learning check Select the correct answer for aluminum: 13 Al From the periodic table is a metal in group 3A. A. Number of valence electrons 1) 1e- 2) 2e- 3) 3e- B. Electron transfer to form an octet 1) loss of 3e- 2) gain of 3e- 3) gain of 5e- C. Ionic charge of aluminum ion 1) Al3- 2) Al+ 3) Al3+ D. The electronic conﬁguration for the aluminum ion 2 2 6 2 1 2 2 6 1) 1s 2s 2p 3s 3p 2) 1s 2s 2p Learning check Select the correct answer for aluminum: 13 Al From the periodic table is a metal in group 3A. A. Number of valence electrons 1) 1e- 2) 2e- 3) 3e- B. Electron transfer to form an octet 1) loss of 3e- 2) gain of 3e- 3) gain of 5e- C. Ionic charge of aluminum ion 1) Al3- 2) Al+ 3) Al3+ D. The electronic conﬁguration for the aluminum ion 2 2 6 2 1 2 2 6 1) 1s 2s 2p 3s 3p 2) 1s 2s 2p Formation of negative ions (anions) In ionic compounds, nonmetals: - Form octets by gaining electrons. - Form negative ions with the electron conﬁguration of the nearest noble gas (in same period). - Negative ions (anions) have more electrons than protons, then have negative ionic charge. 2- Gained e- Ionic charge: - - 1- Group 7A nonmetals → 7 val.e , gain 1 val.e → ion Group 6A nonmetals → 6 val.e-, gain 2 val.e- → ion 2- - - 3- Group 5A nonmetals → 5 val.e , gain 3 val.e → ion - Negative ion of chlorine (Cl ) Chlorine 17Cl - Is a nonmetal located in group 7A. Conﬁg. 1s22s22p63s23p5 - Achieves an octet by gaining (adding) 1 electron to its 7 val. e- - Forms a negative ion with an ionic charge of 1- (Cl1- or Cl- ) - The ion’s new elec. conﬁg. is 1s22s22p63s23p6 (as noble gas 18Ar) Chlorine Chlorine atom: ion: 17p+ 17p+ - - 17e 18e 0 1- Neutral Negative charge charge Learning check Select the correct answer for sulfur: 16 S From the periodic table is a nonmetal in group 6A. A. Number of valence electrons 1) 4e- 2) 6e- 3) 8e- B. Electron transfer to form an octet 1) loss of 2e- 2) gain of 4e- 3) gain of 2e- C. Ionic charge of sulfur ion 1) S2- 2) S2+ 3) S- D. The electronic conﬁguration for the sulfur ion 1) 1s22s22p63s23p6 2) 1s22s22p63s23p4 Learning check Select the correct answer for sulfur: 16 S From the periodic table is a nonmetal in group 6A. A. Number of valence electrons - - - 1) 4e 2) 6e 3) 8e B. Electron transfer to form an octet 1) loss of 2e- 2) gain of 4e- 3) gain of 2e- C. Ionic charge of sulfur ion 2- 2+ - 1) S 2) S 3) S D. The electronic conﬁguration for the sulfur ion 1) 1s22s22p63s23p6 2) 1s22s22p63s23p4 Ionic charges by group (PT) Group numbers of Periodic Table can be used to determine the ionic charges for ions of the representative elements. Elements in group 4A and 8A do not typically form ions. Metals: charge = group num. Non-metals: charge=group num. - 8 Ion becomes as the previous noble gas Ion becomes as the next noble gas Be2+ Ga3+ Se2- Types of bonds A chemical bond is an attraction between atoms that allows the formation of compounds. These bonds can occur through the transfer or just the sharing of their valence electrons → electrons found in the outermost level, allowing the formation of bonds. There are two main types of these: ionic and covalent When atoms or ions come close to each other, electrostatic forces generate interaction. If the resulting attraction forces are strong, a chemical bond forms. Types of chemical bonds Ionic bonds: Involve transference of valence electrons from atoms of a metal to atoms of a nonmetal to form ions. - Loss of electrons by a metal. - Gain of electrons by a nonmetal. Covalent bonds: Involve sharing of valence electrons between the two atoms. (No ions are involved). - Usually form between atoms of two nonmetals, or a nonmetal and a metalloid. (No metals are involved). Ionic bonding Ions with positive charges are attracted to ions with negative charges. Ionic bond is the attractive force between oppositely charged ions that holds them together to form a compound. Na+ → ionic bond ← Cl- It is considered that the more electronegative element attracts the electrons it needs to complete its outermost shell and achieve the conﬁguration of a noble gas. Ionic bonds form between simple ions when metal atoms lose valence electrons and those electrons are gained by nonmetal atoms. Ionic bonding Due to the electrostatic attractions between the elements and the charge each one acquires, this type of compound can conduct electric current. Characteristics of ionic bond: Also called electrovalent bonds Formed between a metal and a non-metal Involves the transfer of electrons from one atom to another Compounds formed have high melting and boiling points Conduct heat and electricity Ionic bonding Another way to identify if a compound has an ionic bond, besides the union between a metal and a non-metal, is by its electronegativity difference on the Pauling scale. An ionic bond has an electronegativity difference greater than or equal to 1.7 (△𝐸≥1.7). Covalent bonding There are elements that need to share their electrons to complete the octet rule; one or more pairs of electrons can be shared. This type of bond is generated when two non-metals unite. The covalent bond results from sharing valence electrons. This type of bond generally occurs when elements are close together in the periodic table and have very similar electronegativities. For example, non-metals overlap or superimpose their orbitals to share the pair of electrons that forms a covalent bond. Covalent bonds The number of covalent bonds that a nonmetal forms is usually equal to the number of electrons it needs to acquire a stable noble gas conﬁguration (to complete the octet). Covalent bonds Atoms share electrons, so they do not form ions. It is formed by the union of two non-metals with equal or very similar electronegativity. Compounds formed have low melting and boiling points. They are not conductors of heat or electrical current. Types of covalent bonds Depending on their electronegativity difference, can be: 1. Polar covalent bond 2. Non-polar covalent bond Polar covalent bond Occurs between two non-metals, but one of them has a higher electronegativity than the other, causing the electron density to be skewed towards one side, resulting in the formation of a dipole. one part of the molecule has a slightly negative charge density and the other part a slightly positive charge density, with the negative part having greater electron density. 𝛿+,𝛿− On the Pauling scale, polar covalent compounds have an electronegativity difference greater than or equal to 0.5 but less than 1.7 (0.5≤△𝐸≤1.7) Water molecule: Electronegativity of oxygen = 3.5 Electronegativity of hydrogen = 2.1 △𝐸=3.5−2.1=1.4 Non-polar covalent bond Occurs between two nonmetals whose electronegativity difference ranges from 0 to 0.4 on the Pauling scale(△𝐸≤0.4) It does not generate dipoles, meaning there are no areas with higher charge density than others. This type of bond occurs between elements that form diatomic molecules where the electronegativity difference is zero or very close to zero. An example is the diatomic molecule of oxygen (O2), in which the electronegativity difference is as follows: Electronegativity of oxygen = 3.5 △𝐸=3.5−3.5=0 The nonpolar molecule O2 looks like this: References Oroz Bretón, N. y Socorro Leránoz, A. B. (2017). Destellos de luz: ( ed.). Pamplona, Spain: Universidad Pública de Navarra. Recuperado de https://0-elibro-net.biblioteca-ils.tec.mx/es/ereader/consorcioitesm/60809?page=101. Dingrando, L. Gregg, K. Hainen, N. Wistrom, (2003). C. Química Materia y Cambio. Mc Graw Hill: México

Module 1: Atomic Structure & Chemical Properties PDF

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