Reference for all bromine oxidation states (-1 to +7) with example compounds, identification rules, and a compound-to-oxidation-state lookup.
Use the Bromine Oxidation States
Atomic #35BrBromine
Atomic Mass
79.904 u
Group
17 (VIIA)
Period
4
Block
p-block
Electronegativity
2.96 (Pauling)
Oxidation States
-1, 0, +1, +3, +5, +7
Bromine has six common oxidation states. Click a state card to see its compounds,
uses, and how to identify it.
Oxidation State −1
The -1 state is by far the most stable and most common oxidation state for bromine. With a 4p5 configuration, bromine needs only one more electron to achieve the stable noble-gas configuration of krypton (4p6). In this state bromine exists as the bromide ion (Br-), occurring in ionic salts and in covalent bonds where bromine is more electronegative than its partner.
How to Identify This State
Assign Br = -1 whenever it bonds to a metal or to hydrogen (in HBr). In ionic bromide salts the metal cation balances the -1 charge. When Br bonds to non-metals less electronegative than itself, it also takes -1.
Example Compounds
Formula
Name
Assignment
Notes
HBr
Hydrogen Bromide
H = +1, molecule neutral → Br = -1
Strong acid in water (hydrobromic acid). Produced by reacting H2 and Br2 or by hydrolysis of PBr3.
NaBr
Sodium Bromide
Na = +1, compound neutral → Br = -1
Ionic salt; used in photography, medicine (sedative), and as a source of Br2.
KBr
Potassium Bromide
K = +1, compound neutral → Br = -1
Used in IR spectroscopy pellets and historically as an anticonvulsant.
MgBr2
Magnesium Bromide
Mg = +2, 2 Br → total -2 → Br = -1
Lewis acid catalyst in organic synthesis; also found in some mineral brines.
CH3Br
Methyl Bromide
C-Br bond: Br more electronegative → Br = -1
Fumigant and methylating agent; now heavily restricted due to ozone-depleting properties.
Common Uses
Ionic bromide salts in photography and medicine
Hydrobromic acid in organic synthesis
Flame retardants (bromide-based compounds)
Drilling fluids (calcium bromide, zinc bromide)
Synthesis of organobromine compounds
Oxidation State 0
Elemental bromine (Br2) is assigned oxidation state zero by convention. It is a reddish-brown liquid at room temperature — the only non-metallic element that is liquid at standard conditions. Both bromine atoms in Br2 are identical, so no electron transfer occurs and both carry oxidation state 0.
How to Identify This State
Any sample of elemental Br2 or a Br2-containing mixture where bromine has not reacted carries oxidation state 0. No calculation is needed — elemental halogens are always 0.
Example Compounds
Formula
Name
Assignment
Notes
Br2
Elemental Bromine
Elemental → Br = 0
Dense reddish-brown liquid; bp 59 °C; pungent, toxic vapour. Produced by oxidizing bromide with Cl2 or by electrolysis of brine.
Common Uses
Production of organobromine flame retardants
Synthesis of pharmaceutical intermediates
Water treatment (bromine tablets for pools and spas)
Production of ethylene dibromide (historically as antiknock additive)
Starting material for all bromine chemistry
Oxidation State +1
Bromine reaches +1 when it bonds to a more electronegative atom and donates one electron formally. The most important example is hypobromous acid (HBrO), formed when Br2 dissolves in water. Hypobromite (BrO-) is the conjugate base and is responsible for bleaching and disinfection in brominated water systems.
How to Identify This State
In HBrO: H = +1, O = -2. Molecule neutral: +1 + Br + (-2) = 0 → Br = +1. In the hypobromite ion BrO-: O = -2, ion has charge -1 → Br + (-2) = -1 → Br = +1.
Example Compounds
Formula
Name
Assignment
Notes
HBrO
Hypobromous Acid
H = +1, O = -2, neutral → Br = +1
Weak acid; formed when Br2 dissolves in water. Potent disinfectant and bleaching agent.
BrO-
Hypobromite Ion
O = -2, charge -1 → Br = +1
Base form of HBrO at high pH. Present in brominated pool water; bleaches and disinfects.
BrF
Bromine Monofluoride
F = -1, neutral → Br = +1
Unstable interhalogen; disproportionates readily to Br2 and BrF3 or BrF5.
BrCl
Bromine Monochloride
Cl = -1, neutral → Br = +1
Interhalogen used as a disinfectant and water treatment agent.
Common Uses
Swimming pool and spa disinfection (hypobromite)
Bleaching in water treatment
Organic synthesis (electrophilic bromination via Br+ transfer)
Interhalogen chemistry research
Oxidation State +3
The +3 state is less common than -1 or +5 but appears in bromous acid (HBrO2) and bromine trifluoride (BrF3). BrF3 is a remarkable solvent for fluorination reactions and acts as both an acid and a base in its own autoionization. Bromous acid is an intermediate in bromate reactions and is not very stable.
How to Identify This State
In HBrO2: H = +1, 2 O = -4, neutral → +1 + Br - 4 = 0 → Br = +3. In BrF3: F = -1 (always), 3 F = -3, neutral → Br = +3.
Example Compounds
Formula
Name
Assignment
Notes
HBrO2
Bromous Acid
H(+1) + 2O(-2) + Br = 0 → Br = +3
Unstable intermediate in bromate/bromite chemistry; short-lived in solution.
BrO2-
Bromite Ion
O(-2)×2, charge -1 → Br = +3
Conjugate base of HBrO2; used in some water treatment processes.
BrF3
Bromine Trifluoride
3F(-1) + Br = 0 → Br = +3
Reactive liquid; auto-ionizes to [BrF2]+ and [BrF4]-; powerful fluorinating agent; T-shaped molecular geometry.
Common Uses
BrF3 as a non-aqueous solvent for fluorinations
Intermediate in bromine oxyacid chemistry
Research into interhalogen compounds
Oxidation State +5
The +5 state is the most practically important of the positive bromine states. Bromic acid (HBrO3) and bromate salts (BrO3-) are strong oxidizing agents widely used in chemistry. Bromine pentafluoride (BrF5) is a potent fluorinating agent. The Belousov-Zhabotinsky oscillating reaction depends critically on bromate chemistry.
How to Identify This State
In HBrO3: H = +1, 3 O = -6, neutral → +1 + Br - 6 = 0 → Br = +5. In BrO3-: 3 O = -6, ion charge -1 → Br - 6 = -1 → Br = +5. In BrF5: 5 F = -5, neutral → Br = +5.
Example Compounds
Formula
Name
Assignment
Notes
HBrO3
Bromic Acid
H(+1) + 3O(-2) + Br = 0 → Br = +5
Strong acid and oxidizing agent; exists only in solution; decomposes on concentration.
KBrO3
Potassium Bromate
K(+1) + 3O(-2) + Br = 0 → Br = +5
Stable crystalline salt; used in food industry (dough improver, now banned in many countries) and as an oxidant in titrations.
BrO3-
Bromate Ion
3O(-2), charge -1 → Br = +5
Central species in the Belousov-Zhabotinsky oscillating reaction; classified as a potential carcinogen in drinking water.
Analytical chemistry (potassium bromate as a primary standard)
Oscillating chemical reactions (Belousov-Zhabotinsky)
Fluorination reactions (BrF5)
Historical use as a flour improver (bromate)
Oxidizing agent in organic synthesis
Oxidation State +7
Bromine's highest oxidation state is +7, reached in perbromic acid (HBrO4) and perbromate salts (BrO4-). The +7 state was long considered inaccessible for bromine — more difficult to achieve than for chlorine (+7 in HClO4) or iodine (+7 in HIO4) — but HBrO4 was first synthesized in 1968 and is a stronger oxidizing agent than perchloric acid.
How to Identify This State
In HBrO4: H = +1, 4 O = -8, neutral → +1 + Br - 8 = 0 → Br = +7. In BrO4-: 4 O = -8, ion charge -1 → Br - 8 = -1 → Br = +7.
Example Compounds
Formula
Name
Assignment
Notes
HBrO4
Perbromic Acid
H(+1) + 4O(-2) + Br = 0 → Br = +7
Strong acid and powerful oxidizing agent; synthesized in 1968; more reactive than perchloric acid.
KBrO4
Potassium Perbromate
K(+1) + 4O(-2) + Br = 0 → Br = +7
Stable salt of perbromic acid; studied for oxidizing properties in analytical chemistry.
BrO4-
Perbromate Ion
4O(-2), charge -1 → Br = +7
Tetrahedral oxyanion; stronger oxidizing agent than perchlorate (ClO4-).
Common Uses
Oxidizing agent in specialized chemical reactions
Research into halogen oxyacid chemistry
Comparison studies with perchlorate and periodate
Synthesis of exotic bromine(VII) compounds
Select a compound from the list to see the oxidation state of bromine
with a step-by-step calculation.
Select a compound to see the oxidation state calculation.
Oxidation state of Br:
Step-by-step
Oxidation State Summary
State
Stability
Key Example
Notes
-1
Most Common
HBr
Bromide ion; gains one electron to reach noble-gas configuration; found in ionic salts and HBr.
0
Elemental
Br2
Assigned by convention to elemental bromine; only non-metallic element liquid at room temperature.
+1
Uncommon
HBrO
Hypobromous acid and hypobromite; formed when Br2 dissolves in water; disinfectant.
+3
Rare
BrF3
Bromine trifluoride and bromous acid; BrF3 is an important fluorinating solvent.
+5
Common
KBrO3
Bromate; strong oxidizing agent; key to Belousov-Zhabotinsky oscillating reactions.
+7
Rare
HBrO4
Perbromate; highest state; stronger oxidizer than perchlorate; first isolated in 1968.
Copied!
Summary
Bromine (Br, atomic number 35) is a halogen in Group 17 that commonly exhibits six oxidation states: -1, 0, +1, +3, +5, and +7. Its ground-state electron configuration is [Ar] 3d10 4s2 4p5 — one electron short of a full p shell, making -1 the most stable and most common state. Unlike fluorine (which is always -1), bromine can reach high positive states by using vacant 4d orbitals to expand its valence shell when bonded to more electronegative atoms such as fluorine or oxygen. This tool lets you explore each oxidation state, see example compounds, and look up the oxidation state of bromine in specific compounds.
How it works
Click an oxidation state card (-1, 0, +1, +3, +5, or +7) to open its detail panel.
The detail panel shows a description, example compounds, identification rules, and common uses.
Use the Compound Lookup tab to select a known bromine compound and see a step-by-step calculation.
Click any formula badge to copy it to your clipboard.
Switch between the Explorer and Compound Lookup tabs using the tab bar.
Use cases
Students revising halogen chemistry and oxidation state rules for exams.
Chemistry teachers preparing reference materials on Group 17 elements.
Researchers checking the oxidation state of bromine in a specific reagent.
Learners working through redox problems involving bromine oxyacids.
Anyone comparing the reactivity of bromine across its oxidation states.
Frequently Asked Questions
Bromine commonly exhibits -1, 0, +1, +3, +5, and +7. The -1 state (bromide) is most common, found in NaBr, HBr, and most ionic bromides. Zero is assigned to elemental Br2. Positive states (+1, +3, +5, +7) occur when bromine bonds to more electronegative atoms — principally fluorine and oxygen in the oxyacids and oxyfluorides.
Bromine has the electron configuration [Ar] 3d10 4s2 4p5 — one electron short of the stable noble-gas configuration. Gaining one electron to become Br- (bromide, 4p6) is energetically favorable and very common. This -1 state appears in virtually all ionic bromide salts and in HBr.
Assign known oxidation states: H = +1 and O = -2 (three O atoms, total -6). The molecule is neutral, so: +1 + Br + 3(-2) = 0 → Br + 1 - 6 = 0 → Br = +5.
Fluorine is always -1. With five F atoms the total is -5. The molecule is neutral: Br + 5(-1) = 0 → Br = +5.
Yes. The +7 state appears in perbromic acid (HBrO4) and perbromate salts. Historically, +7 bromine compounds were harder to synthesize than the chlorine or iodine analogues, but HBrO4 is now well-characterized and is a strong acid and oxidizing agent.
Fluorine is the most electronegative element, so it always attracts bonding electrons toward itself and can only be 0 (elemental F2) or -1. Bromine is less electronegative than both fluorine and oxygen. When bonded to F or O, bromine formally loses electrons, giving it positive oxidation states. Bromine also has accessible 4d orbitals that allow valence-shell expansion to accommodate ten or more electrons.
We use cookies to improve your browsing experience, analyze traffic, and personalize content. By clicking “Accept,” you agree to our use of cookies. Learn more
